Naming Ionic and Molecular Compounds PDF

Naming Ionic and Molecular Compounds What are Compounds? ○ Compounds form when Chemical Bonds are created. Bonds resulting from the force of attraction between atoms … this occurs when elements gain or lose electrons and become “charged”…...

Naming Ionic and Molecular Compounds What are Compounds? ○ Compounds form when Chemical Bonds are created. Bonds resulting from the force of attraction between atoms … this occurs when elements gain or lose electrons and become “charged”… EXPLANATION: If an atom LOSES an electron it becomes positive. The atom that ACCEPTS the electron becomes negative. The atoms are now ions and are attracted to each other. “The ions “stick together” and an ionic bond is now formed. Ions  Properties of Ionic Compounds All solid at room temp NaCl(s) (high melting and boiling points) Retain crystal shape Dissolve in water (majority) NaCl(aq) Always conduct electricity in solution What are Ionic Bonds? An ionic bond forms when one of the 2 elements loses an electron and the other gains an electron to become a stable ion. Normally when a metal (left of step ladder) meets with a non-metal (right of step ladder)  For sodium chloride, there is one sodium ion for Ionic bonds form crystal lattice due to the every chloride ion alternating +/- ion arrangement. (they are in a 1:1 ratio). This is very stable arrangement, so all ionic compounds are solid at room temperature. Strongest bond unless water is around Naming Ionic Compounds In 1911 the International Union of Pure and Applied Chemistry (IUPAC) came up with a naming system. Rules for Type 1 1) The first word is the metal or cation 2) The first word is named exactly like the element. EX. Na is sodium 3) The last word is the nonmetal or anion It ends in an ide ending Ex. Chlorine goes to Chloride Oxygen goes to Oxide Selenium to Selenide Phosphorus to Phosphide  NAMING (Binary) Ionic Compounds IUPAC came up with Formula Cation Anion Name of Compound rules for naming NaCl(s) Na+ Cl- Sodium 1. Name the cation first chloride using the name of the metal BaF2(s) 2. Name the anion K3N(s) (non-metal) 3. Change the ending of the MgO(s) anion to -ide CaBr2(s) Writing Ionic Compounds When writing formulas of binary ionic compounds the symbol for the elements are written in the same order as they appear in the name.  MAKING (Binary) Ionic Compounds ALL compounds are Ex: calcium nitride neutral so the ions must balance out The name does not mention how many because we can use ions to figure it out  Subscript numbers are used to indicate the ratio of the ions in the compound. The charges on the ions must balance in the chemical formula, since ionic compounds are electrically neutral. Identify the ions and their charges. Na+ Cl- 1. Determine the ratio of charges needed to 1:1 balance. 2. The charge on the metal ion crosses to become the Na+1 Cl-1 subscript on the non-metal ion. The charge on the non-metal ion crosses to become the Na1 Cl1 subscript on the metal ion. Reduce the ratio of subscripts in NaCl the formula. Practice Problems 1) aluminum fluoride 2) silver sulfide 3) potassium iodide 4) zinc nitride Workbook 5) calcium oxide page 18 Rules Type 2 (Multivalent/ Binary Compounds) Some metals have more than one choice in the amount of electrons they lose. Ex. Ni2+ or Ni3+, Au3+ or Au+ Note the one on the periodic table that is found on the top is the one most commonly found in nature. Ionic compounds with multivalent elements must have Roman numerals after the name of the positive (metal) ion to indicate the charge on that ion. Compound Name Formula iron (III) chloride lead (IV) oxide nickel (III) sulfide copper (II) fluoride chromium (III) sulfide **Use roman numerals ONLY when the metal element is multivalent.  Naming is the same except for the cation has Roman Numerals Ex. Pb2+ + Cl- becomes PbCl2 named Lead (II) chloride Try the following examples: Cu2+ + O2- Cr3+ + S2- Mn4+ + O2- Workbook page 19 Polyatomic Ions Polyatomic ions consist of a group of atoms combined together that exist as a single unit with an overall electric charge. Most polyatomic ions have a negative charge, which means they behave as non-metals. This means that they are always written last in the formula. The one exception: ammonium ion Examples: Compound Name Formula barium hydroxide iron (III) carbonate copper (I) permanganate gold (III) nitrate ammonium phosphate potassium dichromate  REMEMBER: Use the same general procedure as you did with binary ionic compounds but when you need more than one polyatomic ion … … USE BRACKETS! barium hydroxide ammonium phosphate Try these examples yourselves: Ca(OH)2 K2CrO4 Co(NO2)2 ammonium phosphide magnesium borate barium carbonate ammonium nitrate potassium nitrite Workbook calcium phosphate page 20 Molecular Compounds A molecule is two or more non-metal atoms bonded together. Each molecule is independent of the next and is not part of a lattice. Binary molecular compounds are formed between two non-metal elements.  Properties of Molecular Compounds Not strong enough to take electrons Share pairs of electrons (covalent bond)  3 main types 1. nonmetal and nonmetal 2. hydrogen and metal 3. diatomic (2 of the same element bonded together) or polyatomic (multiple of the same element bonded together) Covalent Bonds – Sharing Electrons! Remember that non-metals need to gain electrons to have a full outer shell. When non-metal atoms combine, the only way this can be achieved is if they share their outer electrons. a pair of shared electrons two chlorine atoms one chlorine molecule  Since electrons are being shared, there is a strong force of attraction between the two atoms. This force is a covalent bond. two pairs of shared electrons an oxygen atom and a water molecule two hydrogen atoms  Example: carbon dioxide e- e- 2e 2e 2e 2e 8p 2e 2e 2e 6p 2e 8p 2e e- e- Carbon Oxygen Oxygen **MEMORIZE THESE** The vast majority of elements exist in nature as single atoms. These are called monoatomic. There are a few diatomic elements (exist as pairs of atoms), WHICH YOU MUST MEMORIZE. “I Bring Clay For Our New House.” **MEMORIZE THESE** There are two polyatomic elements which also must be memorized: “And four Paving stones for eight Steps.” Naming Molecular Compounds Steps for Naming Binary Molecular Compounds In your Data Book!!  If NO prefix on the first name it is assumed to be a one (don’t have to write mono) Ex. CO carbon monoxide dinitrogen monoxide oxygen difluoride sulfur trioxide Try these problems: SO2(g) CS2(g) N2O3(g) CCl4(l) P4O10(s) Workbook pages 21-22 Atomic Bonding Song https://www.youtube.com/watch?v=ljvX-RMv_lw Properties of Ionic and Molecular Compounds Properties of Ionic Compounds 1. Crystalline solids (made of ions) at room temperature 2. High melting and boiling points 3. Will retain their crystal shape when broken 4. Most are soluble in water to some extent 5. When dissolved in water, the solution will conduct electricity. A solution that conducts electricity is called an electrolyte. Solubility of Ionic Compounds Solubility = measure of how well a substance dissolves in a solvent. Ionic compound dissolved in water = aqueous (aq) solution. This compound has a high solubility. Ionic compound not soluble in water = a solid precipitate. This compound has a low solubility. Predicting Solubility To predict whether or not a particular combination of ions will form a soluble compound or not, the solubility table is used. In your data booklet! 1. Locate negative ion at the top of the table. 2. Below this, find the positive ion either in the row labelled “very soluble” or “slightly soluble”. 3. The ions found in the “slightly soluble” will not dissolve in water and will form a precipitate. 4. The ions found in the “very soluble” row will dissolve in water. Examples (write chemical formula and identify solubility): 1. Calcium sulfite 2. Ammonium sulfide 3. Silver sulfate 4. Potassium bromide Properties of Molecular Compounds 1. Solids, liquids, or gases (made up of individual molecules) at room temperature 2. Low melting and boiling points 3. Will crumble easily when broken 4. Only some are soluble in water 5. Molecular compounds dissolved in water do not conduct electricity. Important Molecular Compounds water H2O (l) hydrogen peroxide H2O2 (l) The names, states, and ammonia NH3 (g) formulas for the following molecular compounds sucrose C12H22O11 (s) containing hydrogen must be methane CH4 (g) memorized! propane C3H8 (g) methanol CH3OH (l) ethanol C2H5OH (l) hydrogen sulfide H2S (g) Special Properties of Water Water is a polar molecule. This means that one end of the molecule has a slight positive charge and the other end has a slight negative charge. Water molecules have a strong attraction for one another. Hydrogen bond = Negative end attacted to neighbours Positive end.  Water's polarity means it has several unique properties: 1. High melting and boiling points 2. Large capacity to absorb heat energy without large changes in temperature 3. Solid state (ice) = less dense than liquid state (water) 4. Surface tension - water molecules are more attracted to each other than to other molecules (cohesion) Workbook page 23 Water’s unique properties have many implications for the existence of life on Earth. HEAT SINK  Kahoots 1. Properties of Ionic and Molecular compounds: https://create.kahoot.it/details/d7df99e9-c2cd -422f-8c82-5d9f6766d7b2 2. Solubility of Ionic Compounds: https://create.kahoot.it/details/bb797aa4-4da 6-48df-af36-0623e081932a SUMMARY IONIC VS. MOLECULAR Acids and Bases  Acids Acid: contains hydrogen and dissolves in water forming a solution with pH less than 7. Ex. include vinegar and stomach acid. Properties of acids: 1. Acids have a sour taste 2. Acids DO NOT have a slippery feel 3. Acids react with metals to form hydrogen gas 4. Acids turn litmus paper red 5. Acidic solutions conduct electricity. This is because all acids dissociate (separate into ions) when they dissolve. Bases Base: compound that dissolves in water forming a solution with a pH more than 7. Ex include ammonia and blood. Properties of bases: 1. Bases have a bitter taste 2. Bases have a slippery feel 3. Bases DO NOT react with metals 4. Bases turn litmus paper blue 5. Basic solutions conduct electricity. They also dissociate into ions. Many basic compounds contain the hydroxide ion (OH-) The pH Scale The pH scale is a logarithmic scale. pH means “power of hydrogen” A decrease in one pH value means a 10 times increase in acidity. Naming Acids Acids always considered aqueous and must have an (aq) subscript after acid formula. Acids can be named in two ways: 1) IUPAC system Places the word aqueous in front of acid name, named as if an ionic compound. *This is the system you need to know and be able to apply* 2) Classical System The classical system uses three different rules, depending on what the acid ends in: -ide, -ate, or -ite. *You only need to recognize these as acids* Classical System Rule #1: When the negative ion ends in -ide, the acid begins with the prefix hydro- and the stem of the negative ion is given the ending -ic, in place of -ide. This is followed by the word acid. Rule #2: When the negative ion ends in -ate, the acid name is the stem of the negative ion giving the ending in -ic, in place of -ate. This is followed by the word acid. Rule #3: When the negative ion ends in -ite, the acid name is the stem of the negative ion given the ending -ous, in place of -ite. This is followed by the world acid. These rules are found on your periodic table; DO NOT MEMORIZE! Examples: Naming bases Bases you will be responsible for naming are those that have a metal cation and hydroxide anion (OH-). Name these as if they are any other polyatomic ionic compound. Ex. magnesium hydroxide Mg(OH)2 Naming Compounds: Summary Ask the questions. Is the compound: Workbook page 24-25 Chemical change is a process that involves re-combining atoms and energy flows. Important Examples of Chemical Change https://www.youtube.com/watch?v=jb4CMnT2-ao A chemical change is a result of a chemical reaction, where new substances are formed. *All chemical reactions have the same general form*: reactants products - Rearrangment of atoms and ions to form products. - Products have different properties than the reactants and energy flows into or out of the system. Evidence of Chemical Change 1. Colour change - one or more of the products has a different color than the reactants 2. Formation of a 3. Formation of a gas - if one of the precipitate - if one products is a gas, of the products is bubbles will appear only slightly soluble in water Evidence of Chemical Change 4. Energy change - this is often noticed as heat or light being absorbed or released A chemical reaction that releases energy is exothermic A chemical reaction that absorbs energy is endothermic 5. Change in odour - the products formed may have a different odour, or no odour at all  Energy change in a chemical reaction Sometimes, to illustrate whether a chemical reaction is exothermic or endothermic, the word “energy” is included in the equation. Endothermic reactions, energy is shown as a reactant Ex. photosynthesis: carbon dioxide + water + energy glucose + oxygen  Exothermic reactions, energy is shown as a product Ex. cellular respiration: glucose + oxygen carbon dioxide + water + energy (ATP) Notice that the two reactions are the reverse of one another! Precipitates Precipitates sometimes form when two aqueous solutions are mixed While both reactants were very soluble in water, one of the products is only slightly soluble (YOU MUST CHECK YOUR SOLUBILITY TABLE FOR ALL PRODUCTS!) NaCl(aq) + AgNO3(aq) NaNO3(aq) + AgCl(s) Review: Chemical reactions One or more substances are produced The products (after reaction) have different chemical properties than reactants (before reaction) Reaction could involve: ○ Formation of a gas = production of bubbles ○ Formation of a solid = solution becoming cloudy and forming a precipitate Review: States ELEMENTS: Their state is shown on the periodic table of elements using a shading system COMPOUNDS: Ionic compounds are always solid, unless dissolved in water and are highly soluble, in which case they are aqueous Acids are always aqueous Molecular compounds can be solid, liquid, or gas and will usually be given to you Writing Chemical Equations Chemical Equations Reactants: substances that go into the chemical reaction Products: substances that come out of the chemical reaction Ex. Photosynthesis Chemical Equations Chemical equations can be written in word form or using chemical symbols, but all use the following format: Reactants Products *Important It does not matter in what order the reactants are listed or what order the products are listed, as long as the reactants are on the left and the products are on the right. Word Equations Word equations are the simplest form of writing a reaction equation. Rather than using formulas, word equations use the names of the compounds and elements. methane + oxygen carbon dioxide + water  Skeleton Equations A chemical reaction equation written using chemical formulas. Skeleton equations are not considered to be a finished because it has yet to be balanced. IT MUST HAVE: 1) chemical formula 2) the state Example #1: Solid magnesium metal reacts with hydrochloric acid to produce aqueous magnesium chloride and hydrogen gas. magnesium + hydrochloric acid magnesium chloride + hydrogen Mg(s) + HCl(aq) MgCl (aq) + H (g) 2 2 Example #2: An iron nail is placed in a solution of copper (II) chloride. As a result, small amounts of copper metal are produced in a solution of iron (II) chloride. iron + copper (II) chloride copper + iron (II) chloride Fe(s) + CuCl2(s) Cu(s) + FeCl2(aq) Balancing Equations All equations must always have the same number of each type of atom on both sides of the equation We do this by adding coefficients in front of the chemical formula The coefficient applies to the entire formula ¡2H2O = two water molecules 2H2O = four hydrogens and two oxygens *Important Steps for Balancing Equations: 1. Determine the correct chemical formula for all reactants and products. a. Check for diatomic molecules b. Check for polyatomic ions c. Indicate correct state of compounds (s, l, g, or aq) 2. Balance metals 3. Balance hydrogen 4. Balance oxygen 5. Recount all atoms 6. If every coefficient can reduce, rewrite the whole equation using the simplest ratio of coefficients. Example #1: CH (g) + O (g) CO (g) + H O(g) 4 2 2 2 Elements Number of atoms in reactants Number of atoms in products C 1 1 H 4 2 O 2 1+2=3 The carbons are balanced, but the hydrogen and oxygen are not. Balance the hydrogen by adding a coefficient before the product containing hydrogen. Example #1: CH (g) + O (g) CO (g) + 2H O(g) 4 2 2 2 Elements Number of atoms in reactants Number of atoms in products C 1 1 H 4 2x2=4 O 2 2+2=4 With the addition of the “2” in front of the water, we balance the hydrogen, but also increase the amount of oxygen on the product side. Example #1: CH (g) + 2O (g) CO (g) + 2H O(g) 4 2 2 2 Elements Number of atoms in reactants Number of atoms in products C 1 1 H 4 2x2=4 O 2x2=4 2+2=4 Now the equation is balanced!  Notes 1. Coefficients must be whole numbers. 2. Do not change subscripts in chemical formula. 3. Do not place coefficients between atoms or ions in a formula. 4. Number of polyatomic ions must be the same on both sides of the equation. Practice #1: Practice #2: Practice #3: Practice #4:  Practice #5: Workbook page 27-31 Types of Chemical Reactions Types of Chemical Reactions Despite there being millions of different possible chemical reactions, they follow certain patterns. Most reactions can be categorized into one of five categories: 1. Formation reactions (or synthesis reactions) 2. Decomposition reactions 3. Hydrocarbon reactions 4. Single replacement reactions 5. Double replacement reactions 1. Formation (Synthesis) reaction Two elements combine to form a compound element + element compound A+B AB These reactions are exothermic They occur without the need of added energy 1. Formation (Synthesis) reaction How to recognize these reactions? The only reaction where both reactants are elements, not compounds. Some challenges with this: Watch for polyatomic elements (eg. O2, H2, etc) If the compound produced is ionic, make sure to balance your charges If the compound includes a multivalent metal, assume the most common, unless otherwise specified Examples: 2. Decomposition reaction A compound breaks down into its composite elements compound element + element AB A+B These reactions are endothermic They require energy to occur Ex. electrolysis of water H 2O H2 O2 2. Decomposition reaction How to recognize these reactions? The only reaction where there is only one reactant Some challenges with this: Watch for polyatomic elements (eg. O2, H2, etc) Examples: 3. Hydrocarbon combustion Hydrocarbons are compounds made only of hydrogen and carbon Examples: octane C8H18(l), glucose C6H12O6(s), methane CH4(g) Combustion means “to burn” Like all fuels, hydrocarbons require oxygen in order to burn Combustion reactions are always exothermic Regardless of what hydrocarbon is burning, the products are carbon dioxide and water vapour. 3. Hydrocarbon combustion How to recognize these reactions? The reactants are a hydrocarbon and oxygen In the description of the reactions, it says “combusts” or “burns” Some challenges with this: While these reactions always produce the same two products, they are the hardest equations to balance. REMEMBER: CHO! (carbon, hydrogen, balancing oxygen last) Example: Write the balanced equation for the combustion of methane gas.  4. Simple replacement reaction Element reacts with a compound and produces a new element and a new compound element + compound compound + element A + BC AB + C If the reacting element is a metal, it will replace the metal cation from the compound. Na(s) + KCl(s) NaCl(s) + K(s) If the reacting element is a non-metal, it will replace the non-metal anion from the compound. Br2(l) + MgCl2(s) MgBr2(s) + Br2(l) 4. Simple replacement reaction How to recognize these reactions? The reactants are a compound and an element NOT oxygen Some challenges with this: You will have to balance the charges in the new ionic compound If the new element is a non-metal, watch for polyatomic elements If the reaction occurs in solution, you will need to check the solubility table for the solubility of the new compound Examples: 5. Double replacement reaction Two compounds exchange ions to form two new compounds compound + compound compound + compound AB + CD AD + CB “A” and “C” are both cations (+ ions) so they will always appear first in the formula “B” and “D” are both anions (- ions) so they will always appear second in the formula H2SO4(aq) + 2 NaOH(aq) → Na2SO4(aq) + 2H2O(l) 5. Double replacement reaction How to recognize these reactions? It's the only reaction where both reactants are compounds Some challenges with this: You will be dealing with two new ionic compounds, so you’ll have to balance the charges for both compounds You will also have to check the solubility table for both new compounds Examples:  Practice: Classify the reactions below as (Formation, Decomposition, Hydrocarbon Combustion, Single Rreplacement or Double Replacement) a) C6H12(l) + 9O2(g) 6CO2(g) + 6H2O(g) b) CaCl2(aq) + Na2SO4(aq) CaSO4(s) + 2NaCl(aq) c) 2Fe(s) + O2(g) 2FeO(s) d) Be(s) + 2LiBr(aq) BeBr2(aq) + 2Li(s) Workbook e) MnI4(aq) Mn(s) + 2I2(s) page 32-33 The Mole What is the “mole”? It is a quantity that measures a large number of atoms. The mole is the SI (system internationale) unit for the amount of substance. 1 mole is 6.02x1023 and was named after the person who discovered it: Avogadro hence – Avogadro’s Number = NA https://www.youtube.com/watch?v=TEl4jeETVmg&t=257s What is the mole? It is the number of atoms found in 12 grams of carbon-12 (specific isotope) Carbon-12 was picked so molar mass would be similar to atomic mass Carbon 12 is also easily acquired Why do we use the mole? Atoms are too small to count individually. We are lucky because for balanced equations it is the ratio of the different species that is important. Example 1C6H12O6 + 6O2 6H2O + 6CO2 How does this help me? By using a quantity called molar mass we can quickly count the number of atoms of a given substance. (We we’ll do these calculations later).  Molar Mass Molar mass of a compound = how much one mole of that compound weighs. Unit: g/mol (grams per mole) Formula symbol “M” It is found by adding up the total weight of its parts (elements). Individual (elements’) molar mass is found on the periodic table Atomic Molar Mass The mass of one mole of an element is its atomic molar mass ○ it is measured in grams per mole (g/mol) The atomic molar mass of each element is listed on the periodic table – you’ve already been using it. ○ carbon: 12.01 g/mol ○ sulfur: 32.07 g/mol How do we know the molar mass of compounds? How can you find the molar mass of compounds? To find the molar mass of a compound, you have to: 1. write the chemical formula, 2. list the number of atoms of each element, 3. multiply this number by the molar mass of the element. Molar mass of compounds To calculate the molar mass of a compound add together the atomic molar masses of all the atoms in the compound. Calculate the molar masses of the following compounds: Example: Water H2O(l) H = 1.01 g/mol x 2 = 2.02 g/mol O = 16.00 g/mol x 1 = 16.00 g/mol 18.02 g/mol What this means is that if you counted 6.02x1023 molecules of water it would weigh 18.02g (always 2 decmial spots). Try these ones: Na2SO4(aq) (NH4)2SO4 Na = 22.99 g/mol x 2 = 45.98 g/mol S = 32.07 g/mol x 1 = 32.07 g/mol O = 16.00 g/mol x 4 = 64.00 g/mol 142.05 g/mol Calculating mole and mass In order to properly determine substance quantities to use in a chemistry lab, you must determine the total number of moles or the total mass required of that substance. Once you have determined the molar mass, you know the # of grams in 1 mole of a substance. (ie. Cl2 is 70.90 g/mol) The formula that we use is: Where: n = number of moles (mol) m = mass (g) M = molar mass (g/mol) Can you find this in your data booklet? Manipulating the formula Steps to solving conversion problems: 1. Determine what you are looking for (mass or moles) 2. Calculate the molar mass of the substance 3. Manipulate the formula to solve for the unknown 4. Insert values in correct places 5. Solve. Practice: How many moles of silicon are in a 56.18-g sample? How many moles of potassium fluoride are in a 25.0-g sample? Practice: How many moles of silicon are in a 56.18-g sample? m = 56.18 g MSi = 28.09 g/mol n = m = 56.18 g = 2.00 mol M 28.09 g/mol How many moles of potassium fluoride are in a 25.00-g sample? m = 25.00 g MKF = (1 x 39.10) + (1 x 19.00) = 58.10 g/mol n = m = 25.00 g = 0.43 mol M 58.10 g/mol Manipulating the formula You can also use the formula to find the mass of a sample, if you know: the number of moles in the sample the identity of the sample (this allows you to find molar mass) Using the formula to find mass requires you to rearrange the formula to solve for mass m = nM Example: What is the mass of 10.0 mol of water? n = 10.0 mol MH2O = 18.02 g/mol (calculated in a previous question) m = nM = (10.0 mol)(18.02 g/mol) = 180.2 g Practice: What is the mass of 5.0 mol of NaOH? What is the mass of 4.3 mol of ammonia, NH3? Practice: What is the mass of 5.0 mol of NaOH? n = 5.0 mol MNaOH = (1 x 22.99) + (1 x 16.00) + (1 x 1.01) = 40.00 g/mol m = nM = (5.0 mol)(40.00 g/mol) = 200 g What is the mass of 4.3 mol of ammonia, NH3? n = 4.3 mol MNH3 = (1 x 14.01) + (3 x 1.01) = 17.04 g/mol m = nM = (4.3 mol)(17.04 g/mol) = 73.27 g The “mole” and the Conservation of Mass The Law of Conservation of Mass states that the total mass of the reactants is equal to the mass of the products. In balancing equations, you applied this Law: the coefficient you add in front of a compound in a balanced equation is the number of moles of that substance e.g. In the reaction 4 Na(s) + O2(g) 2Na2O(s) 4 moles of sodium reacted with one mole of oxygen to produce 2 moles of sodium oxide Workbook page 34-35