Physical Sciences 3C Criterion 7 Booklet 2024 PDF

**Physical Sciences 3C** **Criterion 7** **Chemical Structures and Properties** **(Bonding, Precipitation & Organic Chemistry)** **[Will Walker, Eunji Kim, Jason Hoare, Elizabeth College, 2021.]{.smallcaps}** **Chemical structures and properties (Criterion 7)** **This thread describes the properties of atoms that lead to chemical interactions. This knowledge can be used to explain and predict the chemical properties, structures and behaviour of substances.** **In this booklet you will learn more about:** **Chemical properties and the structures of atoms** - **the structure of the periodic table is based on the electron configuration of atoms. Similarities and trends in the observable** - **properties of elements, including chemical behaviour and reactivity, are evident in periods and groups in the periodic table** - **division of elements into metals and non-metals** - **elements are arranged into groups of similar elements with similar properties. Main features, including common ionic charges, of groups 1, 2, 17, 18 (or I, II, VII, VIII)** - **reactivity trends in periods 2 and 3 and groups I, II and VII in the periodic table (qualitative)** - **the properties of atoms, including their ability to form chemical bonds, are explained by their electron configurations.** **Properties and structures of materials** - **the type of bonding within substances explains physical properties** - **the structure and properties of metallic, ionic and covalent substances** - **chemical bonds are caused by electrostatic attractions that arise because of the sharing or transfer of electrons between participating atoms; the valency is a measure of the number of bonds that an atom can form** - **ions are atoms or groups of atoms that are electrically charged due to an imbalance in the number of electrons and protons; ions are represented by formulae that include the number of constituent atoms and the charge of the ion, for example SO~4~^2-^, Na^+^** - **the properties of ionic compounds, for example, high melting point, brittleness, ability to conduct electricity when molten or in solution, are explained by modelling ionic bonding as ions arranged in a crystalline lattice structure with forces of attraction between oppositely charged ions** - **naming and finding the formula of ionic compounds using tables of common anions and cations** - **the characteristic properties of metals, for example, malleability, thermal conductivity, electrical conductivity are explained by modelling metallic bonding as a regular arrangement of positive ions (cations) made stable by electrostatic forces of attraction between these ions and delocalised electrons that are free to move within the structure** - **covalent substances are modelled as molecules or covalent networks that comprise atoms that share electrons, resulting in electrostatic forces of attraction between electrons and the nucleus of more than one atom** - **the distinction between intra and inter molecular forces in covalent molecular elements and compounds i.e. strong forces between atoms and weak forces between molecules.** - **electron dot diagrams for molecules of elements and covalent compounds (e.g. simple hydrocarbons and simple common molecular compounds)** - **focus only on water as an example of a highly polar covalent molecule (only refer to the concept that covalent molecular substances can have differing degrees of polarity)** - **polar molecules (only use water as an example) have increased attraction between molecules resulting in increased:** - **melting points** - **boiling points.** - **naming of covalent molecular compounds based on formulae and vice versa** - **elemental carbon exists as a range of allotropes including graphite, diamond and fullerenes, with significantly different structures and physical properties** - **prediction of the type of structure likely to be present in an element or compound by investigating its physical properties.** - **the presence of specific ions in solutions can be identified using analytical techniques such as flame tests (excludes the need to recall specific colour for elements) or chemical reactions** - **use the solubility table to predict products of precipitation reactions; write overall and net ionic equations for reactions and identify spectator ions** - **H~2~, O~2~ and CO~2~ can be identified using simple gas tests** - **identification of unknown inorganic compounds based on solubility.** **Carbon Compounds** - **carbon forms aliphatic hydrocarbon compounds including alkanes, alkenes, alkynes, cyclic alkanes and cyclic alkenes, with properties that are influenced by the nature of the bonding within the molecules** - **structure and naming of organic compounds using IUPAC nomenclature. The stem to contain a maximum of 10 carbon atoms. This is limited to branched and unbranched alkanes, alkenes, alkynes and cyclic organic compounds containing one or more atoms of F, Cl, Br and I (NO other functional group chemistry)** - **the concept of an isomer, writing structural formulae for a given molecular formula** - **saturated and unsaturated hydrocarbons** - **the distinctions between empirical, molecular and structural formulae (structural formula must include all constituent atoms and bonds)** - **simple reactions of alkanes, alkenes and cyclic organic compounds:** - **complete and incomplete combustion reactions to given products** - **substitution reactions with X~2~ (X =halogen)** - **addition reactions with H~2~, X~2~ and HX** - **test for unsaturation using bromine solution.** **\ ** **ELEMENTS and The Periodic Table: METALS** As you explore the elements in the Periodic Table you will notice that there are more trends than just the trends in atomic radius that you have already seen. You will notice that metals (eg magnesium, aluminium, gold) are on the left hand side of the table, with non-metals (eg nitrogen, oxygen, neon) on the right. Metals and non-metals are separated by the semi-metals (eg silicon). Of the 100 or so elements that occur in nature, approximately 70 are **metals**. **Metals** thus make up the majority of the elements in the Periodic Table. Some metals are found uncombined in nature, i.e. as pure elements: gold, silver and to a lesser extent, copper. It is no surprise then, that ancient civilisations only had access to these metals. The vast majority of metals are found in compounds called minerals. A **mineral** is a naturally occuring compound of a metal found in the earth's crust (the lithosphere). An **ore** is a mixture that contains enough of a metal by weight that it is economical to mine. Ores contain minerals and other materials such as clays. For example, bauxite is the ore from which the mineral alumina (Al~2~O~3~) is obtained, and aluminium is extracted from alumina. As knowledge of the natural world increased over time, humans were able to extract more metals from the naturally occuring minerals in their local environments. The heat from a fire provides sufficient energy to extract copper from some of its minerals, and copper was used by early human civilisations. On the other hand, extracting aluminium requires electricity, a relatively recent discovery. The most abundant metals in the earth's crust are iron and aluminium. A metal's abundance and the costs associated with mining and purification will influence the metal's selling price, along with demand for the metal. Next time you're browsing the internet look up the price per kg of some metals. An important feature of metals is that they can form **homogeneous** **mixtures** with other elements, called **alloys**. Alloys are extremely important in all fields of everyday life, and we make alloys because they have more useful **properties** than the uncombined metal. For example, steel (an alloy of iron) is much stronger than iron alone. Other examples of alloys include: - solder: used to join metals together, particularly in electronics, to join wires, and in plumbing. There are various types of solder, typically they are a mixture of lead and tin, although solders used to join plumbing pipes are *not* allowed to contain lead. - brass: mainly made from copper and zinc. Brass has many uses, inlcuding jewellery, locks, plumbing, and musical instruments. - steel is typically an alloy of iron with a few percent of carbon, and many other elements may be present in small amounts. For example, stainless steel usually contain about 11% chromium. Stainless steel is used to make kitchen sink tops, and it does not rust. Ordinary steel rusts easily. Metals have different **properties** when compared to non-metals. **\ ** **COMPARING METALS & NON-METALS** ================================= A comparison of the physical and chemical properties of metals and non-metals is shown below. +-----------------------------------+-----------------------------------+ | **The Characteristic Properties | **The Characteristic Properties | | of Most Metals** | of Most Non-Metals** | +===================================+===================================+ | 1. Metals have lustre (they are | 1. Non-metals are not lustrous. | | shiny). | | +-----------------------------------+-----------------------------------+ | 2. Metals are solids at room | 2. Some non-metals are solid, | | temperature. (What exception | some are liquid and some are | | to this do you know?) | gases. | +-----------------------------------+-----------------------------------+ | 3. Metals have high density. | 3. Non-metals have low density. | +-----------------------------------+-----------------------------------+ | 4. Metals are good conductors of | 4. Non-metals are thermal | | heat. | insulators. | +-----------------------------------+-----------------------------------+ | 5. Metals are good conductors of | 5. Non-metals are electrical | | electricity. | insulators. | +-----------------------------------+-----------------------------------+ | 6. Metals have high strength. | 6. Solid non-metals have low | | | strength. | +-----------------------------------+-----------------------------------+ | 7. Metals are malleable (can be | 7. Solid non-metals are brittle. | | beaten into shape) and | | | ductile (can be stretched | | | into a wire). | | +-----------------------------------+-----------------------------------+ | 8. Metallic oxides are basic. | 8. Non-metallic oxides are | | | acidic. | +-----------------------------------+-----------------------------------+ | 9. Ions of metals have a | 9. Ions of non-metals have a | | positive charge. | negative charge. | +-----------------------------------+-----------------------------------+ The typical properties of metals listed apply to most metallic elements such as copper, zinc, and gold but even so, there are exceptions. For example, mercury is a liquid at room temperature and metals like sodium and potassium have a low density and are soft enough to be cut with a sharp knife. Your teacher will demonstrate this to you in this course. The non-metallic elements include gases such as oxygen (O~2~), nitrogen (N~2~), helium (He), the liquid bromine (Br~2~) and solids such as iodine, sulfur and phosphorus. More about non-metals later. What about the physical properties of metals? To understand those we need to consider the structure of metals and the bonding between the particles that make up metals. **\ ** **METALLIC BONDING** **Metallic bonding** is the term used to describe the bonding that occurs within a sample of a metal. A metal consists of a 3D lattice (or network) of closely packed metal ions, surrounded by a "sea" of delocalised, mobile valence electrons. The lattice is continuous, extending throughout the metal. This is the **structure** of a typical metal. Have a look at the nearest metallic object. At a particle level you are looking at a 3D lattice of cations surrounded by rapidly moving electrons. This can be visualised as follows: +-----------------------------------+-----------------------------------+ | ![](media/image2.png) | "sea" of delocalised electrons | +===================================+===================================+ | http://friedbiochem.weebly.com/me | | | tals-and-its-physical-properties. | | | html | | +-----------------------------------+-----------------------------------+ In the metallic bonding model, metal ions and electrons are held together by strong, electrostatic attraction between the positively charged cations and negatively charged electrons. This is the **bonding** in a sample of a metal. **Explanation of some of the properties of metals** =================================================== The metallic bonding model can be used to to explain many of the physical properties of metals: - - - - - - **CHEMICAL PROPERTIES OF METALS** ================================= Metals react with a range of chemicals, and you will study chemical reactions in more detail in Criterion 8. One of the common chemical characteristics of metals is that they all tend to react by **losing** electrons and hence the formation of **positive ions or CATIONS.** This process usually involves the metal forming a positive ion with the same number of electrons as its nearest Noble Gas (He, Ne, Ar, Kr, Xe, Rn). You have seen this already in Part 1 of Criterion 4. e.g. Consider the changes in electron configuration for some familiar metals as they undergo reactions: i\) Sodium Na = 2,8,1 Loses its 1 outer (valence) electron to become Na^+^ which is: Na^+^ = 2,8 This is the same electron arrangement as a Ne atom. ii\) Calcium Ca = 2,8,8,2 Loses its 2 outer (valence) electrons to become Ca^2+^ which is: Ca^2+^ = 2,8,8 This is the same electron arrangement as an Ar atom. iii\) Aluminium Al = 2,8,3 Loses its 3 outer (valence) electrons to become Al^3+^ which is: Al^3+^ = 2,8 This is the same electron arrangement as a Ne atom. Despite this general similarity of forming positive ions by electron loss, metals vary tremendously in their tendency to lose electrons. Metals that tend to lose their electrons very readily are highly reactive metals e.g K, Na, Ca. Metals that do not tend to lose their electrons very readily are unreactive metals or 'noble' metals e.g. Ag, Au, Pt. This makes these metals good choices for jewellery... nobody wants their gold jewellery to corrode after a shower or a swim. By carrying out chemical tests in the laboratory later in the course you will probably achieve a listing of chemical reactivity similar to the one shown below. ![](media/image5.png) One of the interesting reactions of metals is the reaction of Group I metals with water. (Group I metals are known as the **alkali metals**). For example, sodium reacts with water to make sodium hydroxide solution and hydrogen gas. 2Na(s) + 2H~2~O(l) 2NaOH(aq) + H~2~(g) Your teacher will demonstrate this and other reactions for you during this course. You will notice that the Group I metals become **more reactive** as you move down the Group. Why is this? **Reactivity of Group I metals -- increasingly reactive moving down the Group** - In Criterion 4 you learned that metals lose their valence electrons to form positively charged ions (cations). When the Group I metals react they lose their single valence electron. Moving down Group I, the atomic radius increases, and thus the valence electron is further and further from the nucleus. - This decreases the electrostatic attraction between the positively charged nucleus and the negatively charged valence electron down the group, making the valence electron easier to remove. - As a result, the Group I metals become more reactive moving down the Group. - Once you understand the atomic radius trend as you move down Group I, and if you relate this to the ease of removing the valence electron, then this reactivity trend is easy to explain. This argument applies to all metals that form cations as we move down the Groups. **ELEMENTS and The Periodic Table: SEMI-METALS** The semi-metals are not dealt with in any depth in this course, but you should be aware of them. They can be found on the roughly diagonal line that separates the metals from the non-metals on the Periodic Table, and are generally considered to be: B, Si, Ge, As, Sb, Te, Po and At. These elements are known as semi-metals (or metalloids) because they have **properties** that are "in-between" those of metals and non-metals. For example, silicon is lustrous and can conduct electricity, but it is brittle rather than malleable. Boron reacts with sodium as a non-metal, but with fluorine as a metal. So they are classified on the basis of their observable **properties**. **\ ** **ELEMENTS and The Periodic Table: NON-METALS** The non-metals are found on the right hand side of the Periodic Table. They can be solids at room temperature (eg sulfur, phosphorus, iodine), liquids at room temperature (eg bromine), or gases (eg oxygen, nitrogen, fluorine, chlorine). You saw typical properties of non-metals in the comparison table above. **Some important groups of non-metals** **The Group VIII elements** are known as the "noble gases" because they are chemically stable, and generally unreactive. This is because they have a full outer electron shell (valence shell). These elements exist as single atoms (ie they are "monatomic"), and as the name of the group suggests, all of these elements are gases under normal conditions. **The Group VII elements** are known as the halogens, and these elements are "diatomic" -- they exist as two atoms bonded together. Fluorine (F~2~) and chlorine (Cl~2~) are gases, bromine (Br~2~) is one of only two liquid elements (the other is the metal mercury, Hg), and iodine (I~2~) is a solid. As you saw in Criterion 4 Part 1, when these elements form ions they do so by gaining one electron to complete their valence electron shell. Unlike the alkali metals, **the halogens become less reactive moving down the Group**. How do we explain this trend? **Reactivity of Group VII elements, the halogens -- decreasingly reactive moving down the Group** - Group VII elements react by gaining an electron to complete their valence shell. - Moving down the Group the atomic radius increases, and thus the valence shell is further from the nucleus. - This causes weaker and weaker electrostatic attraction between the nucleus and an extra electron as we move down the Group. - As a result, the Group VII elements become less reactive moving down the Group. This argument applies to all non-metal atoms that form anions as we move down the Groups. **\ ** **BONDING AND CHEMICAL COMPOUNDS** That brings us to the end of our study of the elements, and the beginning of our study of compounds. Chemical compounds are all around us, and in fact most elements are not found as *elements*, but rather as part of a *compound*. Sodium metal is so reactive that it is never found un-combined; it is always found in compounds. The same is true for nearly all elements. You have already learned that compounds are: - pure substances, containing only one type of particle, - made up of two or more elements chemically bonded together, and - that they have constant composition and properties. (For example, pure water is composed of H~2~O molecules, containing 2 hydrogen atoms and one oxygen atom (always), and it always boils at 100 ^o^C at "sea level" (1 atm of pressure). You already know the names and formulas of some chemical compounds... common examples are: - water: H~2~O - sodium chloride (table salt): NaCl - carbon dioxide: CO~2~ - carbon monoxide: CO - copper(II) sulfate, CuSO~4~ Apart from being found in nature, compounds can also be made in chemical reactions. Before we can learn much more about chemical reactions. we need to be able to **name and write the chemical formula** of compounds. There are two types of compound: **ionic compounds** and **covalent compounds**. We will start our discussion with ionic compounds. We'll learn how to name them, and how to work out their chemical formulas. **IONIC COMPOUNDS** Ionic compounds are made from **ions**. Positively charged ions (cations) attract negatively charged ions (anions), and this electrostatic attraction holds the ions together to form a compound. The names and formulas of the common ions you need to work with in this course are found on the Information Sheet for Physical Sciences 3. **Ionic compounds often contain METAL IONS!** **There are three aspects to writing the names and formulas of ionic compounds:** 1. **simple binary ionic compounds** 2. **compounds containing polyatomic ions** 3. **compounds containing transition metal ions.** **\ ** **1a. NAMING BINARY IONIC COMPOUNDS -- ONE METAL ION AND ONE NON-METAL ION** Write the name of the **metal ion first**, then the name of the **non-metal ion second** (change the end!). For example if sodium reacts with chlorine the compound sodium chloride is formed. Magnesium and oxygen react together to form magnesium oxide. Aluminium and sulfur react together to make aluminium sulfide. Notice the *end* of the name! If I see the formula Al~2~O~3~ I think... aluminium is a metal, so it's an ionic compound.... I'll name the metal ion first, and then the non-metal ion (remembering to change the ending). So, for example: Al~2~O~3~ = aluminium oxide Na~3~N = sodium nitride CaCl~2~ = calcium chloride. The number of ions of each element doesn't affect the name; we **don't** say "calcium dichloride". **1b. WRITING THE FORMULA OF A BINARY IONIC COMPOUND** Compounds have no overall charge; they are neutral. As a result, the total number of positive charges from cations must equal the total number of negative charges from anions. This is what determines the formula of an ionic compound. Some examples to start with... **Sodium chloride:** - sodium atoms form 1+ ions, Na^+^ - chlorine atoms form 1- ions (chloride ions, Cl^-^) - one of the sodium ions and one of the chloride ions will "cancel each other out" - the compound will therefore be neutral overall. - So the formula of sodium chloride is NaCl. This means that the formula unit contains 1 x Na^+^ ion, and 1 x Cl^-^ ion. If you want the equation already: 2Na(s) + Cl~2~(g) 2NaCl(s) **Magnesium chloride:** - magnesium forms Mg^2+^ ions, - chlorine forms Cl^-^ ions (chloride ions), and - we will need 2 of the Cl^-^ ions to cancel the charge on the Mg^2+^ ion. - Therefore magnesium chloride has the formula MgCl~2~. - If you want the equation already: Mg(s) + Cl~2~(g) MgCl~2~(s) One last example... **aluminium sulfide:** - Aluminium forms Al^3+^ ions, and - sulfur forms S^2-^ ions (sulfide ions). - Canceling 3+ and 2- might seem tricky, but what if we do this... - Two of the Al^3+^ ions will give 6+, and three of the S^2-^ ions will give 6-. - As a result the formula of aluminium sulfide is Al~2~S~3~. - The equation for those who are interested: 2Al(s) + 3S (g) Al~2~S~3~(s) **1b cont... AN EASIER WAY - THE "CROSS AND DROP" METHOD...** Let's use aluminium chloride as an example. Simply follow these steps: 1. Write the ions and their charges Al^3+^ Cl^-^ 2. Take the [NUMBER] part of each charge and "cross and drop it" Al^3+^ Cl^-^ Al~1~Cl~3~ 3. We don't write "1s" in chemistry, so now we have... AlCl~3~ 4. Ta dah! That's the formula: AlCl~3~. Let's do another one... What about magnesium nitride: 1. Write the ions and their charges Mg^2+^ N^3-^ 2. Take the [NUMBER] part of each charge and "cross and drop it" Mg^2+^ N^3-^ Mg~3~N~2~ 3. And that's it. That's the formula: Mg~3~N~2~ One more because there's one more rule... calcium oxide: 1. Write the ions and their charges Ca^2+^ O^2-^ 2. Take the [NUMBER] part of each charge and "cross and drop it" Ca^2+^ O^2-^ Ca~2~O~2~ 3. We simplify the ratio, and the formula becomes... CaO This "cross and drop" method works because it results in the same number of positive and negative charges, so the compound is neutral overall. **\ ** **For you to do:** 1\. Complete the table by writing the name and formula of the compound formed by combining the ions in the headings. Follow the example shown. +-----------------+-----------------+-----------------+-----------------+ | | **F^-^** | **O^2-^** | **P^3-^** | +=================+=================+=================+=================+ | **Li^+^** | Lithium | | | | | fluoride | | | | | | | | | | LiF | | | +-----------------+-----------------+-----------------+-----------------+ | **Ca^2+^** | | | | +-----------------+-----------------+-----------------+-----------------+ | **Al^3+^** | | | | +-----------------+-----------------+-----------------+-----------------+ 2\. Write the name and formula of the compound formed when: 3\. Name these compounds: **2. IONIC COMPOUNDS CONTAINING POLYATOMIC IONS -- NAMES AND FORMULAE** Some ions are groups of atoms with a charge. Common examples are: - Sulfate SO~4~^2-^ one sulfur atom, four oxygen atoms, overall charge of 2- - Nitrate NO~3~^-^ one nitrogen atom, 3 oxygen atoms, overall charge of 1- - Ammonium NH~4~^+^ one nitrogen atom, 4 hydrogen atoms, overall charge of 1+ Please refer to the table of common ions on the Information Sheet for Physical Sciences 3 for more example. Ones you should MEMORISE: Sulfate, nitrate, carbonate, phosphate, hydroxide, ammonium... at least. Ask your teacher if you want to know which other ions are useful to remember. When writing the **name** of a compound containing polyatomic ions, follow the same rules as for binary compounds. 1. **Write the name of the metal, or cation first.** 2. **Write the name of the anion second.** Examples: - The compound formed from calcium ions and sulfate ions is calcium sulfate. - The compound formed from ammonium ions and nitrate ions is ammonium nitrate. - The compound formed from ammonium ions and chloride ions is ammonium chloride. - MgCO~3~ is magnesium carbonate. - Li~3~PO~4~ is lithium phosphate. When writing the **formula** of a compound containing a polyatomic ion(s), **use the same "cross and drop" method** **with a slight modification**... For example, what is the formula of magnesium nitrate? 1. Write the ions and their charges Mg^2+^ NO~3~^-^ 2. Take the [NUMBER] part of each charge and "cross and drop it" Mg^2+^ NO~3~^-^ Mg~1~(NO~3~)~2~ 3. We don't write "1s" in chemistry, so now we have Mg(NO~3~)~2~ 4. Notice that we use parentheses ("brackets") to show that we have TWO OF THE WHOLE NITRATE ION in this formula! - Sodium hydroxide is made from Na^+^ and OH^-^ ions. One of each is needed: NaOH - Magnesium hydroxide is made from Mg^2+^ and OH^-^ ions. Two hydroxides needed: Mg(OH)~2~ - Aluminium nitrate is made from Al^3+^ and NO~3~^-^ ions. Three nitrates needed: Al(NO~3~)~3~ - Ammonium sulfate is made from NH~4~^+^ and SO~4~^2-^ ions. Two ammoniums needed: (NH~4~)~2~SO~4~ **\ ** **For you to do:** 1\. Complete the table by writing the name and formula of the compound formed by combining the ions in the headings. Follow the example shown. +-----------------+-----------------+-----------------+-----------------+ | | **OH^-^** | **SO~4~^2-^** | **PO~4~^3-^** | +=================+=================+=================+=================+ | **Li^+^** | Lithium | | | | | hydroxide | | | | | | | | | | LiOH | | | +-----------------+-----------------+-----------------+-----------------+ | **Ca^2+^** | | | | +-----------------+-----------------+-----------------+-----------------+ | **Al^3+^** | | | | +-----------------+-----------------+-----------------+-----------------+ 2\. Write the name and formula of the compound formed when: 3\. Name these compounds: **3. IONIC COMPOUNDS CONTAINING TRANSITION METALS** (this is the last set of rules for ionic compounds!) **Most transition metals can form ions with various charges**. For example, copper can form Cu^+^ and Cu^2+^ ions, iron can form Fe^2+^ and Fe^3+^ ions, etc. As a result, there is no point saying to a chemist "please bring me some iron oxide". They will say... do you mean iron(II) oxide or iron(III) oxide? What's this all about? To make sure that there is no confusion when we name compounds containing transition metals, we include the charge on the ion in Roman numerals in the name. The "(II)" in iron(II) oxide tells me it contains Fe^2+^ ions. Its formula is then.... Fe^2+^ and O^2-^... FeO. The "(III)" in iron(III) oxide tells me it contains Fe^3+^ ions. Its formula is.... Fe^3+^ and O^2-^... Fe~2~O~3~. Notice that **the Roman numeral tells you the CHARGE on the metal ion**, and from that you can work out the formula of the compound. **EXCEPTIONS: We don't do this for zinc or silver. Zinc only forms 2+ ions, and silver only 1+ ions.** **So... ZnO = zinc oxide. AgNO~3~ = silver nitrate.** Some other important metals can form various ions too. We also use Roman numerals to name compounds of tin and lead. SnCl~2~ is tin(II) chloride. SnO~2~ is tin(IV) oxide. PbO is lead(II) oxide. PbCl~4~ is lead(IV) chloride. Make sure you are clear about how these Roman numerals are working and what they mean. **For you to do:** 1\. Complete the table by writing the name and formula of the compound formed by combining the ions in the headings. Follow the example shown. +-----------------+-----------------+-----------------+-----------------+ | | **O^2-^** | **SO~4~^2-^** | **PO~4~^3-^** | +=================+=================+=================+=================+ | **Cu^+^** | Copper(I) oxide | | | | | | | | | | Cu~2~O | | | +-----------------+-----------------+-----------------+-----------------+ | **Cu^2+^** | | | | +-----------------+-----------------+-----------------+-----------------+ | **Cr^3+^** | | | | +-----------------+-----------------+-----------------+-----------------+ 2\. Write the name and formula of the compound formed when: 3\. Name these compounds (hint: you will have to calculate the charge on the transition metal ion first!) **Now practice writing the names and formulas of ionic compounds when they're all mixed up!** ============================================================================================= **Write the name of:** **Write the formula of:** ------------------------ --------------------------- AgCl manganese(II) iodide Ca(OH)~2~ beryllium sulfide LiF aluminium nitrite CuSO~4~ zinc oxide FeS lead(IV) nitrate FeBr~3~ tin(II) carbonate Li~2~O Chromium(VI) oxide **THE STRUCTURE OF IONIC COMPOUNDS** ==================================== **Now you have seen how we write the names and formulas for ionic compounds. Now you are ready to think about the properties of ionic compounds. Surprise, surprise... to do this we will consider their structure and bonding.** ================================================================================================================================================================================================================================= **IONIC COMPOUNDS ARE 3D LATTICES OF ALTERNATING ANIONS AND CATIONS.** **(They are NOT molecules!!)** The structure of ionic compounds can be visualised like this, using NaCl as an example: **The description of the structure above is very important... ionic compounds are 3D lattices of alternating cations and anions. Question: why "alternating"?** =============================================================================================================================================================== **The lattice is as big as the crystal you are looking at! So imagine a crystal of salt from your salt shaker at home... that crystal is one continuous lattice of Na^+^ and Cl^-^ ions. These compounds are therefore *nothing like* individual molecules!** ============================================================================================================================================================================================================================================================= The formula of an ionic compound really only tells us the **ratio of ions in the compound**, not the actual composition of the crystal, which would contain many billions of ions. This type of formula, which tells the ratio rather than the exact composition, is called an **empirical formula.** (Covalent molecules, which do have a specific and fixed composition have what we call a **molecular formula**. More about that later.) **THE BONDING IN IONIC COMPOUNDS** ================================== **IONIC COMPOUNDS CONTAIN STRONG ELECTROSTATIC ATTRACTION BETWEEN IONS OF OPPOSITE CHARGE, THROUGHOUT THE LATTICE.** We can use ionic structure and bonding to explain the properties of ionic compounds: - - - Source: unknown. - **Ionic compounds are often soluble in water.** In other words, many ionic compounds can dissolve in water to produce a solution. This is because the force of attraction between water molecules and the ions is strong enough to pull ions off the lattice and dissolve them. You will learn more about this when you study solutions a little later. For now, think about how common NaCl(aq) is on our planet. - **They do NOT conduct electricity in the solid state.** This is because there are no mobile charged particles to transfer electrical charge. There are no free electrons, and the ions are strongly held within the crystal lattice and cannot move and carry charge. - **They DO conduct electricity in the liquid and aqueous states.** (Liquid state means when melted, or molten. Aqueous state means when dissolved in water to make a solution.) This is because there ARE mobile charged particles to transfer electrical charge in these states. The ions are now free to move and can carry charge. ![molten salt](media/image8.jpeg) Source: unknown **Water of hydration (water of crystallisation)** **Many ionic compounds contain water molecules embedded in the crystal lattice. There is usually a certain number of water molecules per "formula unit". For example, 5 water molecules crystallise with each copper(II) sulfate ion pair in hydrated copper(II) sulfate. This is the blue copper(II) sulfate you are familiar with. We represent this as CuSO~4~.5H~2~O, and we say copper(II) sulfate pentahydrate or copper(II) sulfate-5-water. The water can be removed by heating (physical change), and you will learn about this in Criterion 8.** **For you to do:** 1\. In the following table, identify which species would be electrical conductors and indicate the particles that carry the electrical charge if they do conduct electricity. **CHEMICAL SPECIES** **ELECTRICAL CONDUCTOR?** **MOVING CHARGED PARTICLES** ---------------------- --------------------------- ------------------------------ CuSO~4(aq)~ yes positive & negative ions Ag~(s)~ yes electrons KBr~(s)~ NaI~(l)~ Zn~(l)~ NaOH~(s)~ 2\. Expain, in terms of bonding models, why copper metal is a better electrical conductor than a solution of copper sulphate. **3.** Explain, in terms of the ionic bonding model, why solid NaCl does not conduct electricity but an aqueous solution of NaCl does. **4.** Explain why a piece of copper does not shatter when hit with a hammer whereas a crystal of copper(II) sulfate does shatter. 5\. By considering the ***size*** of the attractive forces between ions, predict which of the two ionic compounds potassium chloride (KCl) and calcium sulfide (CaS) would have the higher melting temperature. (Hint: what are the charges on the ions in the compounds?) Prediction: Reason: **COVALENT COMPOUNDS** Covalent compounds are (usually) made up of non-metals. The non-metals do not form ions -- they **share** electrons rather than lose or gain them, just like the covalent molecular *elements* you saw earlier. Covalent compounds follow different naming conventions compared to ionic substances. The rules to name binary covalent compounds are: - - The second element is given an **-ide** ending. **Prefixes** are used to indicate how many atoms of each element are present in the compound. If there is only one atom of the first element, no prefix is used. If there is only one atom of the second element the prefix is mono-. For example, CO is called carbon monoxide rather than carbon oxide. **The prefixes used to show the number of atoms of each element in the compound are:** **Atom in Molecule** **1** **2** **3** **4** **5** **6** **7** **8** **9** **10** ---------------------- ---------- -------- --------- ----------- ----------- ---------- ----------- ---------- ---------- ---------- **Prefix** **mono** **di** **tri** **Tetra** **penta** **hexa** **hepta** **octa** **nona** **deca** **Question:** Write the name of the compounds shown below. \(a) NO~2~ (b) SO~2~ \(c) PCl~3~ (d) SF~4~ \(e) NI~3~ (f) N~2~O~4~ \(g) P~2~O~5~ (h) N~2~O **Question:** Write the formula of the compounds shown below. \(a) nitrogen monoxide (b) sulfur trioxide \(c) phosphorus trifluoride (d) sulfur hexachloride \(e) nitrogen trichloride (f) dinitrogen pentoxide What will you do if you're not *given* the name or the formula of the covalent compound, but rather have to deduce it for yourself? You'll use the concept of **valency,** or "combining power". Just think of the charge that the elements form on their simple ions, and ignore the positive/negative sign... the combining power, or valence is just the number part. For example, what is the formula of the compound made from phosphorus and hydrogen? - Phosphorus has a valence of 3, and hydrogen has a valency of 1. - Use the "cross and drop" method from before and you get PH~3~. - Or think... "3 hydrogens can combine with one phosphorus"... - Therefore the formula is: PH~3~. Covalent compounds can be much more varied in composition that you might think, so this will only work for simple cases. For example, you saw above in the questions that nitrogen and oxygen can combine to make different compounds in many different ratios. You are not expected to be able to work those out for yourself from scratch -- you'll always be given enough information. Now that you have had some practice writing the names and formulas of covalent compounds, it's time to consider the structure, bonding and properties of covalent compounds. The good news is that you have already seen most of this in the section on covalent molecular *elements*. **SUMMARY AND COMPARISON -- NAMES AND FORMULAS** Writing the **formulas** of compounds if you are given the name. +-----------------------------------+-----------------------------------+ | **If you see a metal or other | **If you only see non-metals in | | cation in the name, use the rules | the name, use the rules for | | for IONIC COMPOUNDS** | COVALENT COMPOUNDS** | +===================================+===================================+ | 1. Write the ions and their | 1\. Simply convert the names to | | charges (cation 1^st^). | symbols. | | | | | 2. Use the "cross and drop" | 2\. Use any prefixes to help | | method. | determine the number of atoms | | | of each element present. | | 3. Write the simplest ratio of | | | ions. | Nitrogen trifluoride = NF~3~ | | | | | 4. Don't write "1s". | | | | | | 5. Use parentheses for multiples | | | of polyatomic ions. | | | | | | Magnesium oxide = MgO | | | | | | Copper(II) nitrate = Cu(NO~3~)~2~ | | +-----------------------------------+-----------------------------------+ Writing the **names** of compounds if you are given the formula. +-----------------------------------+-----------------------------------+ | **If you see a metal or other | **If you only see non-metals in | | cation in the formula, use the | the formula, use the rules for | | rules for IONIC COMPOUNDS** | COVALENT COMPOUNDS** | +===================================+===================================+ | 1\. Write the name of the cation | 1\. Simply convert the element | | first. | symbols to words, changing the | | | end of the last one to end in | | 2\. Write the name of the anion | "ide". | | second (remember to change the | | | ending to "ide" if necessary). | 2\. Use prefexes to indicate the | | | number of each atom present | | 3\. If the metal is a transition | (but don't use "mono" for the | | metal other than zinc or | first element in the name.) | | silver, include its charge in | | | the name. | CO~2~ = carbon dioxide | | | | | 4\. If the metal is tin or lead, | | | include its charge in the name. | | | | | | AlCl~3~ = aluminium chloride | | | | | | Fe~2~O~3~ = iron(III) oxide | | | | | | Mg(NO~3~)~2~ = magnesium nitrate | | +-----------------------------------+-----------------------------------+ **\ ** **COVALENT BONDING** You have already seen that metals lose their valence electrons to form cations, and non-metals can gain valence electrons to form anions. But that's not the whole story. Non-metal atoms can also **share electrons to form what are known as covalent bonds**. Covalent bonds are formed when atoms share electrons. (Metals can also form covalent bonds... do Chemsitry at university to learn about this!). There are two types of covalent structure: **covalent molecules** and **covalent lattices (networks)**. We'll look at covalent molecules first. **COVALENT MOLECULES** You saw above the that the halogens exists as **diatomic** molecules, eg F~2~, Cl~2~, Br~2~, and I~2~. You will already know that oxygen gas exists as O~2~, and you may or may not know that hydrogen gas is H~2~ and nitrogen gas is N~2~. The two atoms in each of these molecules are bonded together through the sharing of electrons -- through **covalent bonding**. The **bonding model of these elements** is known as **covalent molecular**. Each molecule is a separate entity with a fixed composition (eg hydrogen gas is always H~2~). How can we visualise the covalent bonding in these molecules, the sharing of electrons? **a. Sharing 2 electrons (one electron pair, a single bond): hydrogen gas and the halogens** Consider hydrogen gas, the simplest case. It has the formula H~2~ and it is a covalent molecule. Each hydrogen atom has 1 electron, and the first electron shell can hold 2 electrons. Therefore each hydrogen atom in H~2~ can share its electron with the other atom, and both will have two electrons in the valence shell. It is easier to visualise than to read... The two hydrogen atoms are sharing a **single** **electron** **pair**, or two electrons, and this is known as a **single covalent bond**. It allows both atoms to have 2 electrons in the valence shell, now full. The halogens also form diatomic molecules in this way. Consider F~2~. Each fluorine atom has 7 valence electrons, but its valence shell can hold 8. By sharing one extra electron, each fluorine atom can achieve a full valence shell with 8 electrons. (Please note that the inner electron shell has deliberately been left off the diagram below to make the valence electrons clearer.) ![](media/image10.png) **b. Sharing 4 electrons (two electron pairs, a double bond): oxygen gas** Oxygen atoms each have 6 valence electrons, meaning that two more are required for a complete their valence shell with 8 electrons. Thus, each oxygen atom can share two electrons with a second oxygen atom to form O~2~. This sharing of 4 electrons in total, or two pairs, gives what we call a double covalent bond. (Again, only the valence electrons are shown in the diagram below.) **c. Sharing 6 electrons (three electron pairs, a triple bond): nitrogen gas** Each nitrogen atom in N~2~ has 5 valence electrons. Sharing 3 electrons will result in both atoms in N~2~ having a full valence shell, with 8 electrons. (This time all of the electrons are shown.) ![](media/image12.png) These electron-shell diagrams can also be represented using **Lewis electron-dot diagrams.** In these diagrams **[only] valence electrons should be shown**. For example, the electron-dot diagram of an oxygen atom would be: and an oxygen molecule would be **You can see the double bond.** In the spaces below, draw electron-dot diagrams for: Hydrogen gas, H~2~ Chlorine gas, Cl~2~ Bromine liquid, Br~2~ Nitrogen gas, N~2~ -------------------- --------------------- ----------------------- -------------------- **COVALENT MOLECULAR BONDING INVOLVING DIFFERENT ELEMENTS** =========================================================== Sharing electrons, just like gaining/losing electrons, results in a full outer shell, which is a more stable electron arrangement than an incomplete outer shell. Another way of thinking about the formation of covalent compounds is to go back to electron-dot diagrams. **For example, c**onsider the compound formed between the two non-metallic elements phosphorus and chlorine. What will be the chemical formula of the compound? Using the 'electron-dot' method to find the answer... Phosphorus is represented as and chlorine is represented as ![](media/image14.png) These atoms combine by way of sharing ***unpaired*** valence electrons. This means that three chlorine atoms will combine with one phosphorus atom to give PCl~3~. You could have arrived at the same result using the "valency" method. The electron dot diagram is shown below, along with the "structural formula", which shows the bonds. Just like the covalent molecular elements, covalent molecular compounds contain **strong, covalent bonds between atoms, and weak intermolecular forces between molecules.** Another similarity -- the sharing of 2 electrons, or one electron pair, gives a single covalent bond. The three single bonds are shown in the structural formula above, on the right. Note that the "non-bonding electron pair" **is shown in the electron-dot diagram, but is NOT shown in the structural formula.** **DRAWING LEWIS ELECTRON-DOT DIAGRAMS** ======================================= **Electron-dot** diagrams are often referred to as **LEWIS DIAGRAMS** and **are used to represent the sharing of valence electrons in molecular substances. ONLY THE VALENCE SHELL ELECTRONS ARE SHOWN IN THESE DIAGRAMS, FOR EACH ATOM IN THE DIAGRAM.** The Lewis diagrams that you draw will show that all atoms in the molecule achieve a full outer electron shell containing 8 electrons (except for hydrogen whose valence shell is full with 2 electrons). There are some simple rules we can follow for drawing electron-dot diagrams, and the easiest way to go through these is by considering some specific examples. Let's start with one we have already seen, PCl~3~. +-----------------------------------+-----------------------------------+ | 1\. Write the atoms on the page | | | symetrically, making the first | | | atom in the formula the central | | | atom in the diagram. | | +===================================+===================================+ | 2\. Add up ALL the available | Cl has 7 valence electrons and | | valence electrons. | there are 3 Cl atoms 21 electrons | | | | | | Phosphorus has 5 valence | | | electrons. | | | | | | This gives a total of 26 valence | | | electrons. | +-----------------------------------+-----------------------------------+ | 3\. Distribute the valence | ![](media/image16.png) | | electrons *in pairs* around the | | | peripheral atoms (the outside | This has used 24 out of 26 | | atoms), starting with the | electrons... | | bonding pair (the electrons | | | between the central and | | | peripheral atom), so that each | | | of these peripheral atoms has a | | | full valence shell (ie 8 | | | electrons, unless it's hydrogen | | | when it will only have 2 | | | electrons). | | +-----------------------------------+-----------------------------------+ | 4\. Place the remaining pair(s) | | | of electrons on the central | | | atom, as a non-bonding pair. | | +-----------------------------------+-----------------------------------+ | 5\. Now check your diagram: | Yes. | | | | | \- do all atoms have complete | Yes. | | valence shells? | | | | | | \- have you used all of the | | | available electrons? | | +-----------------------------------+-----------------------------------+ NOTE: it is common for different symbols to be used for the electrons of the different elements present. Another example, this time with double bonds, so that we can learn another of the steps... Consider CO~2~... +-----------------------------------+-----------------------------------+ | 1\. Write the atoms on the page | ![](media/image18.png) | | symetrically, with the first | | | atom the central atom. | | +===================================+===================================+ | 2\. Add up ALL the available | Each oxygen has 6 and 6 x 2 = 12 | | valence electrons. | electrons. | | | | | | Carbon has 4. | | | | | | This gives a total of 16 | | | electrons. | +-----------------------------------+-----------------------------------+ | 3\. Distribute the valence | This has used 16 out of 16 | | electrons *in pairs* around the | electrons | | peripheral atoms first, | | | starting with the bonding | AND WE HAVE A PROBLEM... | | electron pair, so that each of | | | these peripheral atoms has a | The carbon atom does not have 8 | | full valence shell. | electrons... | +-----------------------------------+-----------------------------------+ | 4\. We don't have enough | ![](media/image23.png) | | electrons to give carbon a full | | | valence shell! We need to MOVE | | | a non-bonding electron pair on | | | a peripheral atom so that it | | | becomes a shared, bonding | | | electron pair with the central | | | atom. | | | | | | In this example we have to do | | | this twice... | | +-----------------------------------+-----------------------------------+ |... and this gives our electron | and the structural formula would | | dot diagram with two double | be O=C=O | | bonds... | | +-----------------------------------+-----------------------------------+ | 5\. Now check your diagram: | Yes. | | | | | \- do all atoms have complete | Yes. | | valence shells? | | | | | | \- have you used all of the | | | available electrons? | | +-----------------------------------+-----------------------------------+ **For you to do:** In the spaces provided, draw Lewis diagrams for each of the following substances. -------- ------ ------- ------- CCl~4~ F~2~ H~2~O CS~2~ -------- ------ ------- ------- **EXTENSION: Another thing that you can remember if you want to... If the species has an overall charge, take this into account when adding up valence electrons in step 2. For example, NH~4~^+^ has a 1+ charge, so instead of the expected 9 valence electrons, there will only be 8. Simiarly, OH^-^ has a 1- charge, so instead of the expected 7 valence electrons, there will be 8. See if you can use this information to draw the electron dot diagram of NH~4~^+^ now (the ammonium ion).** ===================================================================================================================================================================================================================================================================================================================================================================================================================================================================================================== **FORCES IN COVALENT MOLECULAR BONDING** ======================================== **INTRAMOLECULAR FORCES vs INTERMOLECULAR FORCES** Covalent bonding is **strong!** It results from the electrostatic attraction between positively charged nuclei and the negatively charged shared electrons between them. This bonding is **within** the molecules, and is known as **intramolecular bonding.** These bonds are only broken in **chemical reactions,** not in physical changes such as changes of state**.** The bonding **between molecules, or intermolecular bonding** on the other hand, is **extremely** weak. These intermolecular forces are MUCH MUCH weaker than the electrostatic attraction that occurs throughout a metallic lattice, holding those particles in place. Looking at the temperature required to separate particles (ie the boiling points) of different elements can clarify this... - The higher the boiling point, the more heat energy required to turn the substance from a liquid into a gas, in other words, to separate the particles of the substance from each other. - Therefore, the higher the boiling point the stronger the forces holding the particles in place. BP sodium BP magnesium BP oxygen (O~2~) BP fluorine (F~2~) ----------- -------------- ------------------ -------------------- 880 ^o^C 1100 ^o^C -183 ^o^C -188 ^o^C Because these intermolecular forces are so weak, covalent molecules in the solid state are also **soft**, and covalent molecules are **non-conductors of electricity** (because they have no mobile, charged particles to carry an electric charge). Some examples of elements that are solid covalent molecules are sulfur, which is most commonly found as the yellow solid S~8~ (although we usually refer to it simply as S in chemical equations), and phosphorus, found as P~4~ (and referred to simply as P). They have relatively low MPs and BPs too. So considering hydrogen and the halogens, oxygen and nitrogen allows us to learn about the structure, bonding and properties of covalent molecules. The other type of covalent structure is the covalent lattice (or covalent network). As you learnt by considering covalent molecular elements, within covalent molecules the covalent bonds are strong! These are **not** broken in physical changes such as melting and boiling. You saw that between the molecules there are weak, intermolecular forces. Let's take this idea a little further now... **ALL covalent molecules have weak attractive forces between them (intermolecular forces) called VAN DER WAAL'S FORCES or [DISPERSION FORCES]. Dispersion forces:** **- are the weakest of the intermolecular forces** **- occur between ALL covalent molecules** **- are the only intermolecular forces between non-polar molecules (more about this later)** **- increase in strength with increasing molecular mass** **\ ** **VAN DER WAAL'S (DISPERSION) FORCES** ====================================== Imagine the electrons in a molecule moving around in their shells, between nuclei, non-bonding pairs not between nuclei... VERY BRIEFLY there might be more electrons in one region of the molecule than in another region. This would generate VERY SHORT-LIVED regions of partial positive charge, and partial negative charge on the molecule. These regions can attract neighbouring molecules VERY WEAKLY, due to VERY SHORT LIVED and WEAK electrostatic attraction. **Dispersion forces are extremely weak compared to covalent bonds and thus covalent molecular substances have low melting and boiling points.** The dispersion forces may be so weak that the substance is already a gas at room temperature; e.g. O~2~, CH~4~, F~2~, Cl~2~, N~2~, CO~2~ etc etc. Really engage with this idea of the strength of the forces holding particles together in the solid state by looking up some melting points of the different types of substances you have studied... **Substance** **Type of substance** **Particles present** **Forces b/w particles** **MP (^o^C)** --------------- ----------------------- ----------------------- -------------------------- --------------- NaCl Ionic Cations & anions Electrostatic attraction 801 CH~4~ Mg Si Consider a sample of CF~4~, cooled so much that it is in the solid state. **Within** **the molecule** there are the 4 **strong covalent bonds** (called **INTRAMOLECULAR BONDS**). In a crystal of solid CF~4~ the forces **between** the molecules (intermolecular forces) are weak dispersion forces (these are also referred to as "van der Waal's forces") CF~4~ CF~4~ CF~4~ CF~4~ CF~4~ CF~4~ CF~4~ CF~4~ CF~4~ CF~4~ CF~4~ The dotted lines in the representation above indicate the weak van der Waal's forces (dispersion forces) between molecules of CF~4~. Dispersion forces vary in strength from 'fairly weak' to 'very weak' and increase in strength with increasing molecular mass. The lower the melting point, the weaker the forces. e.g. F~2~ melting point = −220^o^C (extremely weak van der Waal's forces) Cl~2~ melting point = −101^o^C Br~2~ melting point = −7^o^C I~2~ melting point = +114^o^C (fairly weak van der Waal's forces) **\ ** **NON-POLAR vs POLAR MOLECULES -- AN INTRODUCTION ONLY** Think about a Cl~2~ molecule, or an H~2~ molecule, O~2~, or N~2~. Both atoms in these molecules are the same. Therefore, both atoms have nuclei with an equal attraction for electrons. This gives an even distribution of electrons between the nuclei and we say that the N-N bond (for example) is **non-polar.** There are no other atoms or bonds to consider, and so N~2~ molecules are also **non-polar**. For non-polar molecules the only intermolecular forces are dispersion forces. What if the atoms in the molecule are different? What about HCl? A chlorine nucleus has a much stronger attraction for electrons than a hydrogen atom, and so the electrons in the covalent bond are more closely associated with the chlorine nucleus than the hydrogen nucleus. We say that the H-Cl bond is a **polar bond**. ![](media/image25.png) In the diagram the "δ+" and "δ-" symbols indicate regions of partial charge. There is no electron transfer, and no ions are formed. The bond has "more negative" and "more positive" regions, and these are permanent (unlike the short-lived regions of partial charge in non-polar molecules that generate dispersion forces). Perhaps the most familiar and important polar molecule on Earth is **water**. In water, the hydrogen atoms bond at an angle less than 180 degrees, resulting in an oxygen side and a hydrogen side of each molecule. Hydrogen attracts electrons less strongly than oxygen, so water molecules have a partial negative end (the oxygen atom) and a partial positive end (the hydrogen atoms), as shown below. Water molecules in close proximity will tend to align next to each other with each oxygen side facing the hydrogen side of another because of these opposite partial charges. -- -------------------------------------------- ![See the source image](media/image27.png) -- -------------------------------------------- Source: unknown **In addition to dispersion forces, polar molecules also have dipole-dipole interactions between molecules.** (Remember, all molecules have intermolecular dispersion forces.) This is a second type of intermolecular force is stronger than dispersion forces (for molecules of approximately equal size), and so **polar molecules have higher melting and boiling points than non-polar molecules of approximately equal mass.** Water is in fact HIGHLY POLAR, to the extent that its polar bond is given a special name! When a polar bond is EXTREMELY polar, as in water, it is known as a "hydrogen bond". This is beyond the scope of the course, but to give you an idea of the strength of hydrogen bonds (which are still MUCH, MUCH weaker than ionic bonds, metallic bonds and of course covalent bonds), complete the following table. Once you have found the data, reflect on the molecular mass of each molecule (more about this in Criterion 8) and its boiling point... **compound** **classification** **molecular mass** **BP (^o^C)** **compound** **classification** **molecular mass** **BP (^o^C)** --------------------------------------------- -------------------- --------------------------------------------- --------------- -- -------------- -------------------- -------------------- --------------- CH~4~ Non-polar 16 H~2~O VERY polar\* 18 SiH~4~ Non-polar 32 H~2~S polar 34 GeH~4~ Non-polar 77 H~2~Se polar 81 What is the data in this table showing you? What is the data in this table showing you? \*The hydrogen bonding BETWEEN water molecules is so "strong" that it has a higher boiling point than even much heavier, similar compounds. **COVALENT LATTICES (NETWORKS)** Silicon and carbon exist as elements in a different type of covalent structure -- a giant covalent lattice. Rather than having relatively simple, separate molecules, the **whole** crystal is one 'giant' network, or lattice, held together by a continuous linkage of strong covalent bonds. graphite **DIAMOND GRAPHITE** **This network of covalent bonds extending throughout the crystal means each atom is covalently bonded to a number of other atoms, making the substance extremely hard, very strong, with very high melting and boiling points.** **Diamond** and **graphite** are two important covalent network substances that you need to know about. These are two forms of the element carbon and as they have different crystalline structures they are described as **ALLOTROPES** of carbon. Other allotropes of carbon are **fullerenes** and graphene (a single layer of graphite). In diamond, each carbon atom is covalently bonded to four other carbon atoms arranged around it tetrahedrally (see above left). The bonding is continuous in three dimensions and all valence electrons are involved in bonding. This makes diamond extremely hard and a non-conductor of electricity. Diamond is the hardest substance known on earth. Diamond-tipped drills and blades are often used to cut through cement and rock. Diamond's melting point is often quoted as 3550 ^o^C! Graphite is an unusual covalent network substance because it is soft as well as being an electrical conductor. These properties are explained by graphite having carbon atoms bonded in ***layers*** with three valence electrons used to bond each carbon atom to three other carbon atoms (see above right). The fourth valence electron is ***delocalised***, moving between layers. This accounts for graphite's electrical conductivity. Graphite is soft because the individual layers are bonded only weakly to each other. They can slide over each other, making graphite an excellent solid-state lubricant. -------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------- ------------------------ You should also be familiar with another of carbon's allotropes: **fullerenes**. **These are not lattices, but are molecules**, typically containing around 60 carbon atoms, or C~60~. Fullerenes are also known as "buckyballs". Diagram: [www.sciencedirect.com](http://www.sciencedirect.com) ![](media/image29.png) -------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------- ------------------------ **COVALENT COMPOUNDS THAT ARE 3D NETWORKS** The last thing to discuss in this first part of Criterion 7 is the fact that some covalent compounds have a giant 3D lattice (network) structure, and the important example you should remember is silicon dioxide (SiO~2~). In the silicon dioxide covalent lattice every silicon atom is bonded to 4 oxygen atoms (using all 4 silicon valence electrons) and each oxygen atom is bonded to two silicon atoms, using both of oxygen's available valence electrons. **COVALENT NETWORK SUBSTANCES: SUMMARY OF PROPERTIES** +-----------------------------------+-----------------------------------+ | **PROPERTY** | **EXPLANATION** | +===================================+===================================+ | Non-electrical conductors when | There are no ions and all | | solid or molten | electrons are localised in | | | covalent bonds or in the atoms. | | | | | | EXCEPT FOR GRAPHITE in which one | | | electron per carbon atom is | | | delocalised. | +-----------------------------------+-----------------------------------+ | Very high melting points | Very strong covalent bonds extend | | | throughout the crystal lattice. | +-----------------------------------+-----------------------------------+ | Covalent network substances are | All atoms are bound into the | | very hard | crystal lattice by strong | | | covalent bonds. | | | | | | EXCEPT FOR GRAPHITE, in which | | | layers can slide over each other | | | because they are bonded together | | | very weakly. | +-----------------------------------+-----------------------------------+ | Covalent network substances are | The covalent bonding is | | brittle | directional and distortion by | | | force, breaks the covalent bonds. | +-----------------------------------+-----------------------------------+ **SUMMARY AND COMPARISON -- STRUCTURE AND BONDING** +-----------------------+-----------------------+-----------------------+ | **Structure and | **Properties** | **Explanation** | | bonding** | | | +=======================+=======================+=======================+ | **Ionic lattice** -- | \- high melting and | \- electrostatic | | a 3D network or | boiling points | attraction between | | lattice of | | anions and cations is | | alternating cations | \- hard | strong and a large | | and anions, held | | amount of energy is | | together by strong | \- non-conductors of | required to separate | | electrostatic | electricity as solids | them | | attraction throughout | | | | the lattice. | \- electrical | \- ions are fixed in | | | conductors as liquids | place in the lattice | | | | carry charge | | | \- brittle | | | | | \- ions are free to | | | \- many are water | move and carry charge | | | soluble | | | | | \- when the lattice | | | | is deformed, | | | | cation/cation and | | | | anion/anion repulsion | | | | occurs which shatters | | | | the lattice. | | | | | | | | \- interactions | | | | between ions and | | | | water strong enough | | | | to remove ions from | | | | the lattice. | +-----------------------+-----------------------+-----------------------+ | **Metals** -- a 3D | \- good conductors of | \- electrons are | | lattice of metal | electricity as solids | mobile and | | cations surrounded by | and liquids | delocalised and can | | delocalisd, mobile | | move in an electric | | valence electrons. | \- usually high MP | field. | | Strong electrostatic | and BP; strong | | | attraction between | | \- strong | | cations and | \- malleable and | electrostatic | | electrons. | ductile | attraction between | | | | cations and anions | | ![](media/image31.png | | throughout the | | ) | | lattice. Large amount | | | | of heat energy | | [https://quizlet.com/ | | required to separate. | | ](https://quizlet.com | | | | /294622979/metallic-b | | \- when the lattice | | onding-flash-cards/) | | is deformed, the | | | | moving electrons | | | | prevent cation/cation | | | | repulsion. | +-----------------------+-----------------------+-----------------------+ | **Covalent | \- very low MP and | \- weak | | molecules** -- | BP, soft | intermolecular forces | | separate, individual | | between molecules | | and independent | \- non-conductors of | | | particles with fixed | electricity in all | \- no mobile charged | | composition. Strong | states | particles present to | | covalent bonds within | | carry charge | | molecules, very weak | | | | intermolecular forces | | | | (dispersion forces, | | | | dipole-dipole | | | | interactions, | | | | hydrogen bonding) | | | | between molecules. | | | +-----------------------+-----------------------+-----------------------+ | **Covalent lattices** | \- very high MP and | \- all atoms held on | | -- networks or | BP; hard | the lattice by | | lattices of | | strong, covalent | | covalently bonded | \- non-conductors of | bonds; large amount | | atoms. | electricity (except | of energy required to | | | graphite) | remove atoms from the | | May be **3D** eg | | lattice. | | diamond, or **2D** | | | | layered structures eg | | \- all electrons used | | graphite. | | in bonding so no | | | | mobile charged | | ALL atoms covalently | | particles (except | | bonded to other atoms | | graphite which has | | throughout the | | mobile, delocalised | | lattice. | | electrons between | | | | layers) | | ![](media/image33.png | | | | ) | | | +-----------------------+-----------------------+-----------------------+ **Questions for you to do** 1\. Complete the following table: +-----------------------------------+-----------------------------------+ | Write the name of: | Write the formula of: | +===================================+===================================+ | a\. SO~2~ | a\. nitrogen monoxide | | | | | b\. PBr~3~ | b\. sulfur dibromide | | | | | c\. SiCl~4~ | c\. carbon disulfide | | | | | d\. SF~6~ | d\. nitrogen trichloride | +-----------------------------------+-----------------------------------+ 2\. Draw an electron dot diagram of: a\. dihydrogen sulfide b. sulfur difluoride c. nitrogen trichloride 3\. The properties of 6 substances are listed in the table below. Classify each substance as shown in the completed example. Note: (s) = solid; (aq) = aqueous (dissolved in water); NA = not applicable. **Substance** **MP (^o^C)** **Conductivity (s)** **Conductivity (aq)** **Classification** --------------- --------------- ---------------------- ----------------------- -------------------- A 714 ❌ ✔ Ionic compound B 660 ✔ NA C -100 ❌ ❌ D 1170 ❌ ✔ E 3675 ✔ NA 4\. Explain why NaCl conducts electricity when dissolved in water, but SCl~2~ does not. 5. Explain why water has a higher boiling point than oxygen gas, even though oxygen gas is a heavier molecule. **Further review questions and consolidation** A. Complete the following sentences. 1. The majority of the..................................... in the Periodic Table are................................ 2. Metals can form mixtures with other elements called.......................... These are important because of their very useful................................... 3. Metals react with other elements by......................................... electrons to form................................ charged ions, called.............................. 4. The metallic bonding model describes metals as 3D................................. of.............................. surrounded by delocalised, mobile...................................... 5. The bonding between the.............................. and the.................................. is strong.................................. attraction. 6. As a result of the structure and bonding in metals, they have some characteristic properties: a. They are good conductors of electricity because.................................................................. b. They are............................ because cations are tightly packed together on the lattice. c. They generally have.................... MPs and BPs because the............................ between cations and electrons is........................... and extends throughout the lattice. d. They are malleable and................................... because when the lattice is deformed the................................... move to prevent cation-cation..................................., which would otherwise.................................. the lattice, causing the metal to break. 7. Group I metals become................................ reactive moving down the group because moving down the group, the......................................................................... is............................... from the nucleus, and thus the electrostatic attraction between the................................. and this....................... is weaker, and the electron is................................ to remove. 8. Semi-metals such as silicon have some................................. of metals and non-metals. 9. Non-metallic elements are found on the.............................. of the Periodic Table. 10. Some are monatomic, for example the..........................................................., whereas some are diatomic molecules, such as..................,...................,......................... and all of the................................. 11. The atoms in the molecular elements such as chlorine and nitrogen are held together by.......................... bonds, which result from the............................... of electrons. 12. Covalent bonds can be...........................,............................. or............................... depending on how many.............................. are shared. For example, sharing........................... electron pair is a................... bond, and this can be found in............................... molecules. Sharing........................ electron pairs is a double bond, and this can be found in an.............................. molecule. Sharing............................ electron pairs, or............................. electrons, gives a......................... bond, found for example in............................ 13. The covalent bonding in molecules is very strong, and can only be broken in............................... reactions, whereas the..................................... forces,..................................... molecules, are very weak, and these are broken in...................................... changes such as melting and boiling. As a result, the boiling point of covalent molecules is much............................... than the............................... point of metals. 14. Carbon and silicon exist as..................................... lattices. In these structures every atom is bonded to other atoms by strong.................................. bonds. These are very hard to break and as a result, they have extremely................................... MPs and BPs. 15. In....................................... every carbon atom is bonded to four others, leaving no free valence electrons, whereas in........................................ every carbon atom is bonded to only three others, leaving.......................... valence electron per carbon atom unbonded. This electron is delocalised, and moves between layers, allowing................................. to............................................................. 16. Graphite is also a very good..................................... because the layers can.............................. freely over each other. B. Draw an electron dot diagram (Lewis diagrams) of a(n): fluorine atom fluorine molecule oxygen atom oxygen molecule nitrogen atom nitrogen molecule --------------- ------------------- ------------- ----------------- --------------- ------------------- C. Answer the following questions. 1\. Explain why the reactivity of Group I elements increases down the Group whereas the reactivity of Group VII elements decreases down the Group. 2\. Elements A and X are both hard solids at room temperature. A conducts electricity whereas X does not. Identify the type of bonding present in A and X, giving reasons for your answer. 3\. Why do the halogens exist as diatomic molecules containing a single covalent bond? **\ ** **[Precipitation Reactions]{.smallcaps}** **[Introduction]{.smallcaps}** Ionic compounds typically involve metal ions (positive ions or cations) in chemical combination with non-metal ions (negative ions or anions). e.g. AgNO~3~, CaCl~2~, Na~2~CO~3~, Al~2~(SO~4~)~3~, CuSO~4~, NiBr~2~, PbS, Ni(NO~3~)~2~.................. When ionic crystals do dissolve in water they break up into separate ions that then become dispersed throughout the whole aqueous solution. For example, solid NaCl~(s)~ consists of an orderly and close packing of sodium ions (Na^+^) and chloride ions (Cl^−^) into a regular array called a 'crystal lattice'. The ions are held strongly within the lattice by electrostatic attractions between the oppositely charged ions. These attractive forces are called ionic bonds. The dissolving of NaCl~(s)~ in water breaks up the ionic lattice and the ions disperse throughout the whole aqueous solution. A solution of sodium chloride (NaCl~(aq)~) consists of individual ions Na^+^~(aq)~ and Cl^−^~(aq)~ moving independently and randomly throughout the solution. In this particular unit of work we focus on the properties of ionic compounds and specifically whether these compounds will dissolve in water or whether they will not dissolve in water. **Compounds that DO dissolve in water** are described as **SOLUBLE** in water. **Compounds that DO NOT dissolve in water** are described as **INSOLUBLE** in water. Although many ionic compounds are described as 'insoluble' it is true to say that even insoluble ones do dissolve to some small extent. However, for our studies, we can consider that the amount an 'insoluble' compound does dissolve to be effectively ZERO. **[Solubility Table]{.smallcaps}** As a result of experimental investigations, chemists have systematically established which ionic compounds are **soluble** and which ionic compounds are **insoluble**. The results for common ionic compounds are summarized in the table on your Formula Sheet and shown here. **Useful Solubility Rules:** The information in the table allows us to see two important general "solubility rules": - all sodium, potassium, ammonium and lithium compounds are **soluble** - all nitrate and acetate (ethanoate) compounds are **soluble** You will have access to the solubility table in all tests and examinations so there's no need to memorise the information it contains although these 'rules' are worth remembering when it comes to making decisions about the likely formation of precipitates. **[\ ]{.smallcaps}** **[Soluble or Insoluble]{.smallcaps}** You need to be able to use the solubility table on your formula sheet to determine which ionic compounds are **soluble** in water and which ionic compounds are **insoluble** in water. **Example:** Complete the table below by giving the appropriate chemical formula and then writing 'soluble' or 'insoluble' in the third column. **IONIC COMPOUND** **CHEMICAL FORMULA** **SOLUBLE OR INSOLUBLE IN WATER?** --------------------- ---------------------- ------------------------------------ barium chloride aluminium phosphate calcium iodide silver chloride potassium carbonate iron(III)hydroxide **[Mixing Ionic Solutions]{.smallcaps}** When aqueous solutions of magnesium chloride (MgCl~2~) and potassium nitrate (KNO~3~) are mixed, the resulting solution contains Mg^2+^, Cl^−^, K^+^ and NO~3~^−^ ions all moving randomly throughout the whole volume of the solution. This mixed solution is identical to one obtained by mixing equivalent amounts of magnesium nitrate and potassium chloride. Both solutions contain Mg^2+^, Cl^−^, K^+^ and NO~3~^−^ ions all moving randomly throughout the solution. The reason for this is that all the possible ionic compounds which could form from these ions are soluble, that is, magnesium chloride, magnesium nitrate, potassium chloride and potassium nitrate are soluble. a. What ions are present in this mixed solution? b. Which two other ionic compounds could be mixed so as to produce a combination involving the same ions? **[Forming a Precipitate]{.smallcaps}** **Sometimes when clear solutions of two ionic compounds are mixed, they react to produce a solid.** This solid is usually finely divided and slowly settles to the bottom of the test-tube or beaker. **The solid that forms as a result of this reaction and then settles to the bottom is called a PRECIPITATE**. **The reaction that produced the solid is called a PRECIPITATION REACTION.** (To precipitate literally means 'to fall out'.) When solutions of two ***soluble*** ionic compounds are mixed, a precipitate will form if one type of positive ion present can combine with one type of negative ion present to form an ***insoluble*** substance. The **reaction that takes place to form a precipitate is known as a precipitation reaction.** **Example:** When aqueous solutions of barium chloride (BaCl~2(aq)~) and potassium sulfate (K~2~SO~4(aq)~) are mixed, a reaction occurs and a thick white precipitate forms. Using the solubility table, it can be seen that the precipitate is barium sulfate (BaSO~4(s)~) because barium sulfate is **insoluble** in water (but potassium chloride is soluble). This means that whenever a barium ion (Ba^2+^~(aq)~) collides with a sulfate ion (SO~4~^2−^~(aq)~) they join up (bonding ionically) to form the solid BaSO~4~. The net ionic equation is the preferred way of presenting precipitation reactions. **Example:** Using the solubility table, what are two possible solutions you could mix in order to form a precipitate of: \(i) nickel(II) hydroxide? (ii) silver bromide? **Example:** Does a precipitation reaction occur when aqueous solutions of sodium iodide and silver nitrate are mixed? If so identify the precipitate, write the total chemical equation, total ionic equation, net ionic equation and name the spectator ions. **Example:** Does a precipitation reaction occur when aqueous solutions of copper(II)sulfate and sodium sulfide are mixed? If so identify the precipitate, write the total chemical equation, total ionic equation, net ionic equation and name the spectator ions. **Example:** When aqueous solutions of potassium sulfate and barium nitrate are mixed, a white precipitate of barium sulfate forms. Write the net ionic equation and name the spectator ions. **Example:** Water has been accidentally contaminated with lead(II) ions as a result of a spillage of a soluble lead salt into the water. Lead compounds are known to be highly toxic and this contaminated water needs immediate treatment. How could the lead(II) ions be removed from the contaminated water? Explain your procedure and use a chemical equation to supplement your answer. **\ ** **[Tests for gases produced]{.smallcaps}** **Hydrogen gas (H~2~)** The pop-test is used to test for the presence of hydrogen. Hydrogen is a colorless gas that is less dense than air and is highly flammable. The pop-test can be used to identify hydrogen gas. A test tube mixture of hydrogen and air will pop when ignited producing water as a by-product. The chemical reaction is 2H~2~ + O~2~ → 2H~2~O + Energy **Carbon dioxide (CO~2~)** Carbon dioxide passed into limewater (Ca(OH)~2~) gives a milky solution. This is due to the insoluble suspension of [calcium carbonate](https://en.wikipedia.org/wiki/Calcium_carbonate) formed: If excess CO~2~ is added, the following reaction takes place: The milkiness disappears since calcium bicarbonate is water-soluble. **Oxygen (O~2~)** A glowing wooden splint relights in a test tube of oxygen. **[Identification of Unknown Ionic Compounds]{.smallcaps}** Using knowledge of solubility chart for ionic salts, precipitation reactions and common reactions of acids and bases, it is possible to identify an 'unknown' ionic compound using a series of simple tests. An additional test that can be performed is the flame test. **[Flame Test]{.smallcaps}** Scientists are often faced with the problem of identifying what compounds are present in a sample. One method that can be used to identify the metal ions present is the **flame test**. A simple (qualitative) flame test involves dissolving the compound in a flammable liquid such as methanol and igniting. The colour of the flame is then used to determine what cations are present. (An alternate test is to use a non-reactive metal (usually platinum) wire dipped in a paste containing the substance to be tested. This is then heated in a flame and the colour observed.) **Some common metals and their flame test colours** **Metal** **Flame test colour** --------------- ----------------------- **barium** **pale green** **calcium** **yellow -- red** **copper** **green -- blue** **sodium** **orange** **potassium** **lilac** **lithium** **red** More analytical and quantitative tests use controlled and much higher temperatures and special instruments to determine the spectrum of the emission. They can accurately determine what ions are present, and in what relative quantities. This is called Flame Emission Spectroscopy. **Example:** The labels have fallen off three reagent bottles known to contain solutions of sodium nitrate NaNO~3(aq)~, calcium iodide CaI~2(aq)~ and lithium carbonate Li~2~CO~3\ (aq)~. Assuming you had access to solutions of silver nitrate, AgNO~3~, and potassium sulfate, K~2~SO~4~, describe how you would go about identifying each of the solutions correctly. In each case, complete the table with all precititates and spectator ions, where applicable. Describe any **observations** and write **balanced equations** for any reaction that occurs. What additional test could you perform to confirm your tests above. NaNO~3~ CaI~2~ Li~2~CO~3~ ----------- --------- -------- ------------ AgNO~3~ K~2~SO~4~ **\ ** **ORGANIC CHEMISTRY** The term 'organic' chemistry has its origins in the fact that living systems are based on carbon compounds. However, our coverage of organic chemistry will encompass many compounds that are not associated with living things. The principal issue is that the topic of organic chemistry is based on compounds of ***carbon***. The other elements that are often bonded with carbon in these organic compounds are hydrogen, oxygen, nitrogen, sulfur, halogens,............ There are hundreds of thousands of known chemical compounds that are based on carbon structures and tens of thousands of new carbon compounds are being discovered each year. The answer comes down to carbon's bonding capacity. As already discussed in our chemical bonding unit, each carbon atom is able to form four strong covalent bonds and this gives each carbon atom the capacity to form bonds with other carbon atoms. Consequently carbon can form strong linkages between atoms giving rise to numerous compounds involving: - Straight chains of carbon atoms. - Branched chains of carbon atoms. - Rings of carbon atoms. - Combinations of rings and chains of carbon atoms. Most other elements, apart from carbon, are unable to form 'multiple linkages'. As the nu

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