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Main Groups Main Groups Chemistry 223 Chem Dr. Ali Alsalme Modern Atomic Theory  John Dalton 1766-1844 British chemist His Theory:  All substances are made of atoms that cannot be created, divided,...

Main Groups Main Groups Chemistry 223 Chem Dr. Ali Alsalme Modern Atomic Theory  John Dalton 1766-1844 British chemist His Theory:  All substances are made of atoms that cannot be created, divided, or destroyed. Atoms join with other atoms to make new substances. Solid Sphere Model Atoms of the same element are exactly or alike, and atoms of different elements Bowling Ball Model are different in mass and size.  J.J. Thomson 1856-1940 English chemist and physicist; discovered 1st subatomic particles His Theory:  Atoms contain negatively charged particles called electrons and positively charged matter. Created a model to describe the atom as a sphere filled with positive matter with negative particles mixed in. Plum Pudding Model or Raisin Bun Model Referred to it as the plum pudding model.  Ernest Rutherford 1871-1937 New Zealand physicist discovered the nucleus. His Theory:  Small, dense, positively charged particle present in nucleus called a proton. Nuclear Model Electrons travel around the nucleus, but their exact places cannot be described. Neils Bohr 1913 Danish physicist; discovered energy levels His Theory:  Electrons travel around the nucleus in definite paths and fixed distances. Electrons can jump from one level to another level.  Erwin Shrodinger 1924 Austrian physicist; developed the electron cloud model His Theory:  The exact path of electrons cannot be predicted. The region referred to as the electron cloud, is an area where electrons can likely be found. Electron Cloud Model James Chadwick 1932 English physicist; discovered neutrons His Theory:  Neutrons have no electrical charge Neutrons have a mass nearly equal to the mass of a proton Unit of measurement for subatomic particles is the atomic mass unit (amu).  Modern Theory of the Atom Atoms are composed of three main subatomic particles: the electron, proton, and neutron. Most of the mass of the atom is concentrated in the nucleus of the atom. The protons and neutrons are located within the nucleus, while the electrons exist outside of the nucleus. In stable atoms, the number of protons is equal to the number of electrons.  The type of atom is determined by the number of protons. The number of protons in an atom is equal to the atomic number. The sum of the number of protons and neutrons in a particular atom is called the atomic mass. Valence electrons are the outermost electrons. The Periodic Table Elements are organized on the table  according to their atomic number, usually found near the top of the square. Atomic Number This refers to how many protons in the atom. No two elements, have the same number of protons. Atomic Mass Atomic Mass refers to the This is a helium “weight” of the atom. atom. Its atomic mass is 4 (protons It is derived at by adding the plus neutrons). number of protons with the number What is its atomic of neutrons. number? Atomic Mass and Isotopes Isotopes are members of a family of an element that all have the same number of protons but different numbers of neutrons. The number of protons in a nucleus determines the element's atomic number on the Periodic Table. An atomic mass number with a decimal is the total of the number of protons plus the number of neutrons. Valence Electrons Valence electrons are the electrons in the outer energy level of an atom. These are the electrons that are transferred or shared when atoms bond together.  Periodic Table The periodic table organizes the elements in a particular way. A great deal of information about an element can be gathered from its position in the period table. For example, you can predict with reasonably good accuracy the physical and chemical properties of the element. You can also predict what other elements a particular element will react with chemically. Understanding the organization and plan of the periodic table will help you obtain basic information about each of the 118 known elements.  Families Columns of elements are called groups or families. Elements in each family have similar but not identical properties. For example, lithium (Li), sodium (Na), potassium (K), and other members of family IA are all soft, white, shiny metals. All elements in a family have the same number of valence electrons. The groups or families are divided into eight “A” groups and eight “B” groups. The “A” groups are often referred to as the main group or representative elements because they possess a wide range of chemical and physical properties. The “B” groups are referred to as the transition elements. Notice their numbering is not in consecutive order from left to right as with the “A” groups. A more recent numbering system of the groups numbers them consecutively 1-18 from left to right. Periods Each horizontal row of elements is called a period. The elements in a period are not alike in properties. In fact, the properties change greatly across even given row. The first element in a period is always an extremely active solid. The last element in a period, is always an inactive gas. Three Classes of the Elements: Metals, Nonmetals, Metalloids (Semimetals) Nonmetals occupy the upper-right- hand corner. Metalloids (semimetals) are located along the boundary between metals and nonmetals (stair-step line) The majority of the elements are metals. They occupy the entire left side and center of the periodic table. Each of these classes has characteristic chemical and physical properties, so by knowing whether an element is a metal, nonmetal, or metalloid, you can make predictions about its behavior.  Metals All metals except mercury(circled) are solids at room temperature, in fact, most have extremely high melting points. The periodic table shows that most of the metals are not main group or representative elements. They are transition and inner transition metals The elements in Groups 3 through 12 of the periodic table are called the transition elements. All transition elements are metals. Nonmetals Although most elements in the periodic table are metals, many nonmetals are abundant in nature. The nonmetals oxygen and nitrogen make up 99 percent of Earth’ s atmosphere.  Carbon, another nonmetal, is found in more compounds than all the other elements combined. The many compounds of carbon, nitrogen, and oxygen are important in a wide variety of applications. Most nonmetals don’ t conduct electricity, are much poorer conductors of heat than metals. Many are gases at room temperature. Their melting points tend to be lower than those of metals. With the exception of carbon, nonmetals have five, six, seven, or eight valence electrons.  Metalloids (Semimetals) have some chemical and physical properties of metals and other properties of nonmetals. In the periodic table, the metalloids lie along the border between metals and nonmetals. Silicon (Si) is probably the most well-known metalloid. Some metalloids such as silicon, germanium (Ge), and arsenic (As) are semiconductors. The ability of a semiconductor to conduct an electrical current can be increased by adding a small amount of certain other elements. A semiconductor is an element that does not conduct electricity as much as a metal but does conduct slightly better than a nonmetal. Periodic Table Trends Periodic Trends Periodic Trends Atomic Properties (Atomic Radii) Atomic Radii: Half the distance between the centers of neighboring atoms in a solid or a homonuclear molecule. Atoms get smaller from left to right in a period. Atoms get larger from top to bottom in a group When atoms share the same highest energy level, there is a stronger pull on the electrons when there are more protons. Errors in Predicting Atomic Radius Rearrange the following atoms in the order of decreasing atomic radius: Question: Na, Mg, Sr, Cl, S, O, F, Br, Ga, Ti, Rn, Kr Errors in Predicting Atomic Radius Rearrange the following atoms in the order of decreasing atomic radius: Question: Na, Mg, Sr, Cl, S, O, F, Br, Ga, Ti, Rn, Kr Answer: Ti , Rn, Sr, Ga, Br, Kr, Na, Mg, S, Cl, O, F But according to actual atomic radii data, the order should be this: Sr, Na, Ti , Mg, Rn, Ga, Br, Kr, S, Cl, O, F Sr (2.15 Å), Na (1.86 Å), Ti (1.71 Å), Mg (1.60 Å), Rn (1.4 Å), Ga (1.22 Å), Br (1.14 Å), Kr (1.09 Å), S (1.04 Å), Cl (0.99 Å), O (0.66 Å), F (0.64 Å) When comparing periods, it is not true that all atoms in a lower period are larger than the atoms in the period above it. Alkali and alkali earth metals are usually significantly larger than even the largest non- metals Ion Radius It can also be predicted the radius of elements in their ionic form. How do metals and non-metals differ when it comes to comparing their ionic radius to their atomic radius? Metals lose electrons to form ions, so their ions have smaller radii. Non-metals gain electrons to form ions, so their ions have larger radii. Atomic Properties (First Ionization Energy) First Ionization Energy: The minimum energy required to remove the first electron from the ground state of a gaseous atom, molecule, or ion. The energy it takes to remove the outermost electron in an atom (expressed as kJ/mol) If the electronegativity is high, is it harder or easier to remove an electron? i.e. Does it require more or less energy? An atom’ s ability to attract electrons increases as it gets closer to achieving a full octet An atom’ s ability to attract electrons increases when its nucleus is closer to the outermost electrons Based only on their position in the periodic table, arrange the elements in order of increasing ionization energy Li, Br, Zn, La, Si, P, Ga, Cl, Y, Cs Based only on their position in the periodic table, arrange the elements in order of increasing ionization energy ANSWER: Cs, Li, La, Y, Zn, Ga, Si, P, Br, Cl Atomic Properties (Electron Affinity) Electron Affinity (Eea): The energy released when an electron is added to a gas-phase atom. The closer an atom is to achieving an octet, the more stable it will become when an electron is added. The resulting increase in stability causes a release of energy. Hence electron affinity increases from left to right across a period. The distance of the valence electrons from the nucleus increases going down a periodic table. This results in less attraction for the valence electrons, as well as any extra electrons that may be added. The decreased attraction results in less energy released. Thus, electron affinity increases going up a group in the periodic table. Which element in each of the following pairs will have the lower electron affinity? Explain your answer in each case K or Ca )a O or Li )b S or Se )c Cs or F )d Atomic Properties (Electronegativity) Electronegativity (χ): The ability of an atom to attract electrons to itself when it is part of a compound. - Which atom has a greater ability to attract electrons? The closer an atom is to achieving a full octet, the greater its ability to attract electrons. Electronegativity is the measure of an atom’ s ability to attract electrons. Which atom should have the highest electronegativity? Which is second highest? Electrons are closer to the Electrons are farther from the nucleus = stronger pull on nucleus = weaker pull on electrons electrons - Based only on their position in the periodic table, arrange the elements in each set-in order of increasing electronegativity. Li, Br, Zn, La, Si )a P, Ga, Cl, Y, Cs )b Li, Si, Zn, Br, La )a Cl, P, Ga, Y, Cs )b  Bonding Trends (Number of Bonds) Most main group elements form the same number of bonds as the oxidation number. Elements in period three and higher have access to the empty orbitals and can use them to expand their valence shells past the usual octet of e- and therefore do not always follow this rule. Example Typical Number of Typical Oxidation Number of Element Bonds Formed Number Valence e- NaCl(s) 1 1+ 1 Na H2O(l) 2 2 - 6 O HF(l) 1 1 - 7 F Bonding Trends (Size) The smaller the size of the atom, the fewer of the other atoms that can bond with it. In general, only period 2 elements form multiple bonds with themselves or other elements in the same period because only they are small enough for their p orbitals to have substantial π overlap. Hydrogen  Hydrogen In modern periodic table it is located separately Electron configuration is 1s1(similar to the electron configurations of group 1 elements) Classified as a nonmetal. Therefore, it doesn’ t fit into any group. Properties: Colorless, Odorless, Tasteless Similarity to alkali metals: 1. Electronic configuration H = 1s 1 1 2 2 6 1 11Na = 1s , 2s , 2p , 3s K = 1s2, 2s2, 2p6, 3s2 , 3p6, 4s1 19 Electropositive character: H+, Na+, K+ etc..2 3. Oxidation state: +1 4. Combination with electronegative elements: form binary compounds with electronegative elements like alkali metals. Halides: HCl NaCl, KCl etc. Sulphides: H2S Na2S, K2S etc. Similarity to halogens: Electronic configuration: Both contain one electron less than the nearest noble.1 gas configuration 1 1H = 1s (near to 2He) F = 1s2, 2s2, 2p5 (near to Ne) 9 8 2 2 6 2 5 17Cl = 1s , 2s , 2p , 3s 3p (near to 18Ar) 2. Non-metallic character: like halogens, hydrogen is non-metallic in nature. 3. Atomicity: Diatomic molecules. 4. Formation of similar types of compounds: i. Halides: CCl4, SiCl4, GeCl4 ii. Hydrides: CH4, SiH4, GeH4 5. Oxidation state: Na+1H-1 Na+1Cl-1 Difference from alkali metals: Ionization enthalpy:- the ionization enthalpy of hydrogen is very high in.1 comparison to alkali metals. Non-metallic character: alkali metals are typical metals while hydrogen is.2 non-metal. Atomicity: hydrogen is diatomic while alkali metals are monoatomic..3 Nature of compounds: the compounds of hydrogen are predominantly.4 covalent while those of alkali metals are ionic. For example: HCl is covalent while NaCl is ionic. The oxides of alkali metals are basic while hydrogen oxide is neutral. Difference from halogens: Less tendency for hydride formation: Hydrogen has less tendency to take.1 up electron to form hydride ion (H-) as compared to the halogens which from halide ions (X-) very easily. Absence of unshared pairs of electrons..2 Nature of oxides: The oxides of halogens are acidic while hydrogen oxide.3 is neutral. Occurrence of Hydrogen:.4 * Hydrogen, the most abundant element in the universe and the third most abundant on the surface of the globe, is being visualized as the major future source of energy. The Element The elemental form of H is H2 H2 is small and nonpolar so the H atoms can only attract each other through weak London forces. Physical Properties of Hydrogen Density boiling point Melting point Abundance Molar mass symbol Name Z (g.L-1 ) (ºC) (ºC) )%( (g.mol-1 ) 0.089 - 253 (20 K) - 259 (14 K) 99.98 1.008 H hydrogen 1 0.18 - 249 (24 K) - 254 (19 K) 0.02 2.014 2H or D deuterium 1 0.27 - 248 (25 K) - 252 (21 K) radioactive 3.016 3H or T tritium 1 Preparation: Methods for commercial production of dihydrogen Electrolysis of water.1 2H2O(l) 2H 2(g) + O2(g) I t is called green hydrogen. The hydrogen prepared by this method is very high purity. However, this method is not commonly used because it is very expensive. 2. By the reaction of steam on coke :- C + H2O(g) CO + H2 Since the mixture of CO and H2 is used for the synthesis of methanol and several hydrocarbons, it is also called synthesis gas or syngas. 3. Most commercial H2(g) is obtained as a by product of petroleum refining: This method called Steam reforming: CH4(g) + H2O(g) CO(g) + 3H2(g) Can’ t Separate CO from H2 CO(g) + H2O(g) / CO2(g) + H2(g) Membranes that can separate CO2 from H2 Hydrogen produced by steam reforming is termed 'grey hydrogen' when the waste CO or CO2 is released to the atmosphere and 'blue hydrogen' when the CO or CO2 is (mostly) captured and stored geologically. Hydrogen as a fuel source: Light (low density). Clean Burning. Plenty of abundant H in H2O. Properties of Hydrogen: Physical Properties:-  It is slightly soluble in water (about 2 %)..1 It is highly combustible and therefore should be handled carefully..2 It is lightest substance. The weight of one liter hydrogen at STP is only 0.0899 g..3 Chemical properties:  Not very reactive due to high bond dissociation energy (435.88 kJ mol-1at 298.2 K) Combustion: It burns with pale blue flame.1 2H2(g) + O2(g) 2H2O(l) 2. Reaction with metals: Reactive metals like Na, K, Ca, Li and form hydrides. Ca + H2 CaH2 3. Reaction with metal oxides: Hydrogen reduces oxides of less active metals to corresponding metal.  Hydrogen Compounds Hydrogen can form both cations (H +) and anions (H-). Hydrogen has an intermediate electronegativity. Forms covalent bonds with both nonmetals and metalloids. Uses of Hydrogen: Hydrogenation of oils:.1 Vegetable oils are polyunsaturated in nature. The C=C bonds in oils can easily undergo oxidation and the oil becomes rancid i.e., unpleasant in taste. Hydrogenation reduces the number of double bonds. 2. As a reducing agent. 3. In the manufacture of ammonia, metal hydrides, methanol, fertilizers such as urea etc. 4. In the manufacture of synthetic petrol. 5. In the fuel cell for generating electrical energy. Hydrides: Under certain conditions H 2 combines with almost all the elements ,except noble gases to form compounds called hydrides. Hydride s Covalent (Molecular) Ionic (Saline) Hydrides Metallic Hydrides Hydrides Electron-deficie Electron-prec Electron-rich nt ise Three different classifications of binary hydrides:  Covalent (Molecular) Hydrides Ionic (Saline) Hydrides Metallic Hydrides Molecular Hydrides Molecular hydrides are formed when nonmetals form covalent bonds with  hydrogen. Example: HF, HCl, Saline HBr Hydrides: Saline hydrides are formed by the members of the block when they are heated in the presence of H2 Example: 2K(s) + H2(g) ∆ 2KH(s) Metallic Hydrides Metallic hydrides are formed by heating certain block metals in the presence of H2(g) Example: 2Cu(s) + H2(g) ∆ 2CuH(s) Electron deficient: The compounds in which the central atom has an incomplete octet are called electron-deficient hydrides. These types of compounds are normally formed by group 13 elements. Diborane (B2H6) is an example of an electron-deficient hydride. These hydrides do not have a sufficient number of electrons to form normal covalent bonds. Example BH3 and AlH3.These compounds generally exist in polymeric forms like B2H6, B4H10 to make up for the deficiency of electrons. Electron precise: An electron-precise hydride is a type of hydride which has the exact number of electrons required to form normal covalent bonds, no excess no deficient. These are formed generally by group 14 elements like Si and C which have 4 valence electrons. For example, (CH4, SiH4, GeH4, SnH4, PbH4 etc.).  Electron rich hydrides: The hydrides which have excess electrons as required to form normal covalent bonds is called electron rich hydride. For example, hydrides of group 15 to 17 (NH3, PH3, H2O, H2S, H2Se, H2Te, HF etc.) ydrides - YouTube Which of the following hybrids is electron-precise, Electron-deficient, or Electron-rich hydride? B2H6 , NH3 , H2O, CH4 , AlH3 Boron has 3 electrons in the valence shell, and it requires 5 electrons to complete its octet. From the structure we can see that diborane has a 3c-2e bond, hence it is electron deficient. N atom has 5 valence electrons, out of which 3 electrons form covalent bonds with H-atoms and two electrons remain unbonded. All electrons are not used for bond formation. It has an electron-rich center. Oxygen atom has 6 valence electrons of which two electrons form a bond with two hydrogen atoms, but 4 electrons remain unbonded. It has an electron-rich center. Carbon atoms have 4 valence electrons. It forms four bonds with four hydrogen atoms. All electrons are used for bond formation. Therefore, CH4 is electron-precise hydride. Aluminum has 3 electrons in the valence shell, and it requires 5 electrons to complete its octet , if it forms a bond with 3 hydrogen it still requires more electrons to complete its octet. Hence it is electron deficient. Hydrogen Bonding Between H and highly electronegative atoms (ex: N, O, and F). 5% as strong as covalent bonds (between the same atoms). Coulombic interactions between the partially positive charge on a hydrogen atom and the partially negative charge of another atom form the H bond. δ- δ+ O --- H -O  Group 1 The Alkali Metals Group 1: The Alkali Metals The elements in group 1, on the left of the periodic table, are called the alkali metals. The alkali metals all have one electron in their outer shell. lithium These metals are all very reactive and 2,1 are rarely found in nature in their elemental form. They can readily lose the outer shell sodium electron to form positive ions with a +1 2,8,1 charge and a full outer shell. Properties of Alkali Metals  Properties of Alkali Metals Alkali metals are soft enough to be cut with a knife. Softness increases going down the group. Group 1 metals are very reactive with oxygen and must be kept away from oxygen in order to not get oxidized. The speed reacting with oxygen increases going down the group. These alkali metals rapidly react with oxygen to produce several different ionic oxides. Oxides: O2- , peroxides: O22- , super oxide: O -. The usual oxide, M2O, can be formed with alkali metals generally by limiting the 2 supply of oxygen. With excess oxygen, the alkali metals can form peroxides, M2O2, or super oxides, MO2. The elements in group 1 also react with water and form alkaline compounds. This is why they are called alkali metals. They have low melting and boiling points. Generally, the melting point of the alkali metals decreases down the group. This is because as the atoms get larger the distance between the bonding electrons and the positive nucleus gets larger and reduces the overall attraction between them. Hence, the bonds are weaker, and less energy is required to break them. For similar reasons, the electronegativity decreases. All the alkali metals react vigorously with water and increases going down the group. It is an exothermic reaction as it releases a lot of heat. 2Na(s) + 2H2O(l) 2NaOH(aq) + H2(g) The alkali metals are all too easily oxidized to be found in their free state in nature. Great reducing agents. Lithium (Li) Lithium differs slightly from the other element in the group Lithium is harder than other alkali metals. Melting and boiling point is higher than other alkali metals. Out of all the other alkali metals, it is the least reactive metal. It is a strong reducing agent compared to other alkali metal. It is the only alkali metal that forms its monoxide. It is the only alkali metal that forms a nitride. Lithium (Li) Uses: manufacturing of batteries. glass and ceramic industry. polymers and drug industries.  Diagonal relationships The first member of a group is often a bit different. We typically find that its chemistry is more like that of the chemistry of the second member of the next main group than like that of its own group. Diagonal relationships Example: Li and Mg Li is in many ways more like Mg than like the other alkali metals: Li is the only alkali metal to form a nitride, whereas this is a typical reaction for the alkaline earth metals. 6 Li + N2 → 2 Li3N 3 Mg + N2 → 2 Mg3N2 A number of lithium compounds are much less soluble in water than the corresponding compounds of the other alkali metals (e.g. lithium phosphate). Both Li and Mg combine with O2 to form monoxide while other members of their respective groups form peroxide and superoxide. 4 Li + O2 → 2 Li2O 2 Mg + O2 → 2 MgO Li2CO3 decomposes to Li2O and CO2 like the alkaline earth metal carbonates, but the other alkali metal carbonates are thermally stable. Li2CO3 → Li2O + CO2 MgCO3 → MgO + CO2 Compounds of Sodium NaCl (Sodium Chloride commonly known as table salt) Properties: White Solid salt Methods of Obtaining NaCl Mined as rock salt which is a deposit of sodium chloride left as ancient oceans evaporated Obtained from the evaporation of brine (sea water) NaOH (Sodium Hydroxide commonly known as lye) Properties: Uses: Soft.- Inexpensive starting material to produce other sodium salts waxy White - to produce soaps, paper, products that explode, dyes, and Corrosive solid petroleum products. Sodium Hydrogen Carbonate) NaHCO3( (Sodium Bicarbonate) commonly called baking soda. How baking soda works: The hydrogen carbonate reacts with a weak acid (HA) that is present in the batter (sour milk, buttermilk, lemon juice, vinegar.....) HCO3-(aq) + HA (aq) A-(g) + H 2O(l) + CO2(g) The CO2(g) produced causes the batter to rise. Baking powder contains a solid weak acid as well as the sodium hydrogen carbonate therefore CO2(g) is released when water is added. Compounds of Potassium - expensive than Na compounds. - similar properties to Na compounds. Potassium chloride (KCl) KCl is used as a fertilizer, in medicine, in scientific applications, domestic water softeners (as a substitute for sodium chloride salt), and in food processing, where it may be known as E number additive E508. Common Reactions Reaction with Halogens 2M + X2 2MX X2 is any group 7A molecule Reactions with Oxygen 4Li + O2 2Li2O Need excess Oxygen Na + O2 Na2O2 (80%) + Na2O (20%) unbalanced equation Na2O2 + O2 Na2O M + O2 MO2 M = K, Rb, or Cs (Metal superoxide) Reaction with H 2M + H2 2MH Reaction with N 6Li + N2 2Li3N Li only Reaction with Water 2M + 2H2O 2MOH + H2  Group 2 The Alkaline Earth Metals Group 2: The Alkali Earth Metals Electron configuration is ns2+. Classified as metals. These metals are all very reactive and are rarely found in nature in their elemental form. Usually found as +2 cations. Properties of Alkali Earth Metals Beryllium (Be) Beryllium and its compounds are extremely toxic. Used as windows for x-ray tubes: Beryllium is a metal that has low density and low atomic mass, and hence very low absorption of X-rays, making beryllium the preferred choice for X-ray tube windows where low energy transmission is desired. Beryllium is used in gears particularly in the aviation industry. Other beryllium alloys (especially with copper, aluminum, magnesium, or nickel) are used as structural materials for high-speed aircraft, missiles, spacecraft and communication satellites. https://www.youtube.com/watch?v=qy8JyQShZRA&t= 216s Beryllium (Be) The Be atom in the BeCl2 act as Lewis acids and accept electrons pairs form the Cl atoms of the neighboring BeCl2 groups forming a chain of tetrahedral BeCl4 units in the solid. Magnesium (Mg) Occurs in sea water as the mineral dolomite CaCO3·MgCO3 Mg is present in the chlorophyll molecule. Protective oxide forms which protects Mg from extensive oxidation from air. Used as an alloying agent in the manufacturing of airplanes due to it is one-third less dense than aluminum. Magnesium - Periodic Table of Videos – YouTube Magnesium Hydroxide (Mg(OH)2) Commonly called Milk of Magnesia. Because Mg(OH)2 is relatively insoluble in water. It is not absorbed into the stomach and stays in the stomach a long time to react with whatever acid is present. When Mg(OH)2 neutralizes stomach acid it produces MgCl2 which is a laxative therefore Mg(OH)2 should be used reasonably. Magnesium Sulfate (MgSO4) Commonly called Epsom Salts. Is another laxative. The magnesium ions inhibit the absorption of water in the intestines thereby increasing the flow of water through the intestines. Chlorophyll Captures light from the sun and channels its energy into photosynthesis. The role of the Mg2+ ion is to keep the ring rigid thereby ensuring that the photon is not lost as heat before the chemical reaction. Calcium (Ca) Also found in sea water. Ca is the element of rigidity and construction (bones, shells, concrete, mortar, limestone (buildings)...) Obtained by electrolysis or by reduction with aluminum in a version of the thermite process (same for strontium and barium) 3CaO(s) + 2Al(s) ∆ Al2O3(s) + 3Ca(s) Calcium - Periodic Table of Videos – YouTube Compound of Calcium Calcium Carbonate (CaCO3) Most common calcium compound. Occurs naturally in chalk and limestone. CaCO3(s) decomposes into CaO(s) (lime or quicklime) when heated CaCO3(s) ∆ CaO(s) + CO2(g) Compound of Calcium Calcium Oxide (CaO) CaO is called quicklime because the reaction with water is fast and extremely exothermic. CaO(s) + H2O(l) Ca2+(aq) + 2OH-(g) Uses of CaO: Used in iron making. CaO is a Lewis base and reacts with silica in the iron ore to transform it into liquid slag CaO(s) + SiO2(s) ∆ CaSiO3(l) About 50 kg of lime is needed to produce 1 ton of iron Uses of Ca(OH)2: Used as an inexpensive base in industry. To adjust the pH of soils in agriculture. Ca(OH)2 is known as slacked lime because the thirst of lime for water has been quenched or slacked Common Reactions Reaction with Halogens M + X2 MX2 X2 is any group 17 molecule Reaction with Oxygen 2M + O2 2MO Reaction with N 3M + N2 M3N2 High temperatures Reaction with Water M + 2H2O M(OH)2 + H2 Reaction with Ions (acids) M + 2H+ M2+ + H2 Diagonal relationships Example: Be and Al We will see that Be behaves a lot like Al:  Both form a tough and chemically resistant passivation layer. Both have amphoteric oxides..Contrast: basic oxides that define the alkaline earth metals Their oxides don’ t react with water. Contrast: reaction of alkaline earth metals with water to make alkaline.earth hydroxides Both form covalent halides..Contrast: ionic halides of alkaline earth metals  Group 13 The Boron Family Group 13: The Boron Family Electron configuration is ns2np1. Boron exhibits mostly nonmetallic behavior and is classified as a semimetal, whereas the other members of Group 13 are metals. Boron and Aluminum almost always have an oxidation number of +3. The heavier elements of the group are more likely to keep their electrons and can have oxidation numbers of +1 or +3. Properties of Group 13 elements But even the metals have no simple pattern in melting points, although their boiling points do show a decreasing trend as the mass of the elements increases. The reason for this lack of order is that each element in the group is organized a different way in the solid phase. It is in Group 13 that we encounter elements possessing more than one oxidation state. Aluminum has the +3 oxidation state, whether the bonding is ionic or covalent. However, Gallium, Indium, and Thallium have a second oxidation state of +1. For Gallium and Indium, the +3 state predominates, whereas the +1 state is most common for Thallium. Boron B Boron is the only element in Group 13 that is not classified as a metal. It is classified as a semimetal. The element can be obtained from its oxide by heating with a reactive metal such as Magnesium: B2O3(s) + 3Mg( ) 2B(s) + 3MgO(s) Boron - Periodic Table of Videos – YouTube  Boron B High ionization energy. Forms covalent bonds. Small atomic radius. Electron-pair arrangement in diborane, B2H6 Boron has only three valence electrons, so any boron compound that has simple covalent bonding will be electron deficient with respect to the octet rule. Thus, we saw that the simplest boron hydride dimerizes to give B2H6, in which there are two hydridic bridge bonds Geometry of the diborane molecule Elemental boron exists in variety of different structures, one of the more common ones being B12. Because of the three-dimensional network formed by the bonds, boron is very hard and when incorporated in plastics, forms a material that is stiffer than steel yet lighter than aluminum. About 35 percent of boron production is used in the manufacture of borosilicate glass. Compounds of Boron Boric Acid (H3BO3) Properties: White Solid Toxic to bacteria, insects, and humans. Melts at 171º Used as a mild antiseptic and pesticide. C Lewis Acid Also, used in many industrial application. Aluminum Al Most abundant metallic element in the Earth’ s crust. Low density. Strong metal. Amphoteric. Excellent electrical conductor. Commercial source of aluminum is bauxite (Al2O3·xH2O where x ranges from 1 to 3) Aluminium (version 1) - Periodic Table of Videos – YouTube Aluminium (or Aluminum) - Periodic Table of Videos - Compounds of Aluminum Aluminum oxide (Al2O3) Commonly know as alumina. Variety of crystal structures. 90% of the Al2O3 produced is consumed to produce Al. 10% remaining used in a wide variety of applications such as: Fillers: Being chemically inert and white, Al2O3 is a favored filler for plastics - and a common ingredient in sunscreen. Catalysis, Glass, Gas purification (remove water from gas streams), Paints. - Common Reactions Reaction with Halogens 2M +3X2 2MX3 X2 is group 17 molecule, Tl gives TlX as well but no TlI3 Reactions with O 4M + 3O2 2M2O3 Reactions with N 2M +N2 2MN  Group 14 The Carbon Family Group 14: The Carbon Family Electron configuration is ns2np2. The half-filled orbital allows this group to straddle the line between metal and nonmetal. The elements show increasing metallic character as you go down a group. The heavier elements of the group are more likely to keep their s electrons and can have oxidation numbers of + 2 or + 4. Melting and boiling points of the Group 14 elements: Carbon C Central element to life. Carbon has nonmetallic properties. Forms covalent bonds with nonmetals and ionic bonds with metals. Small radius allows for the wide occurrence of C=C and C=O bonds in compounds. Carbon is the only member of group 14 that commonly forms multiple bonds with itself. Carbon - Periodic Table of Videos - YouTube Carbon forms: Diamond sp3 Hybridized carbon (tetrahedral). Only C-C σ bonds. Graphite & Diamond Uses In: Properties: - Jewellery. - Stone polishing and cutting. Rigid Transparent Electrically insulating Solid Good conductor of heat  Carbon forms: Properties: Black Graphite Lustrous Electrically conductive Slippery sp2 hybridized carbon in a hexagonal network Uses In: Electrons are free to move Electrical conductors in industry from one carbon to another Electrodes in electrochemical cells through π network formed Lubricants by the overlap of Lead in pencils unhybridized -orbits on each of the carbon atoms https://www.youtube.com/watch?v=cAjh0lXxt E8  Carbon forms: Soot and Carbon Black Contains very small crystals of graphite https://www.youtube.com/watch?v=FBzmfPZ5Pxo Made: Heating gaseous hydrocarbons near 1000oC in the absence of air Uses In: Reinforcing rubber, pigment, and printing ink  Carbon forms: Activated Carbon (Activated Charcoal) Contains granules of microcrystalline carbon. Made: Heating waste organic matter in the absence of air and then processing it to increase the porosity, producing a very high specific.surface area Uses In: Air purifiers, gas masks, aquarium water filters, water purification plants (remove organic compounds from drinking water). Carbon Compounds Oxides of Carbon Carbon dioxide (CO2) Formed when organic matter burns in a plenty source of air and when animals exhale. CO2 is always present in air, but the burning of fossil fuels is increasing the amount of CO2 in the air which is then in-turn leading to global warming. Uses In: Plants convert carbon dioxide to oxygen during a process called photosynthesis, using both carbon and oxygen to make carbohydrates. It is used as a fire extinguisher.  It is used in promoting the growth of plants in greenhouses. It is used in carbonated soft drinks to make them fizzy. Large quantities of solid carbon dioxide ( in the form of dry ice) are used in large-scale refrigeration. The carbon dioxide released by baking powder or yeast that makes cake batter rise is the best example of the use of carbon dioxide in everyday life. Carbon monoxide (CO) Properties: Formed when carbon burns in a limited source Colorless of air. Odorless This happens in cigarettes and badly tuned Flammable automobile engines Almost Insoluble Toxic Gas CO is a reducing agent and is used in the production of different metals, most notably iron in blast furnaces. https://www.youtube.com/watch?v=otVFDo9YSM8 Silicon Si Central element to electronic technology and artificial intelligences. Larger atomic size than C which results in relatively few compounds that have Si=Si and Si=O bonds. Si compounds can act as Lewis acids whereas C compounds typically cannot. Si compounds can expand its valence shell by using its electrons thereby allowing for the accommodation of lone pair electrons of a Lewis base. Silicon - Periodic Table of Videos - YouTube Second most abundant atom in the earth’ s crust. Occurs widely in rocks as silicates (compounds containing the silicate ion, SiO32-) Pure silicon is obtained by reduction of quartzite (a granular form of quartz) with high purity carbon in an electric arc furnace. https://www.youtube.com/watch?v=HKQ2GaXFI3w SiO2(s) + 2C(s) ∆ Si(s) + 2CO(g) (crude is exposed to Cl2(g)) SiCl4(l) + 2H2(g) Si(s) + 4HCl(g) (purer form of element) Further purification is necessary before silicon can be used in the semiconductor industry Silicon Compounds Properties: Oxides of Silicon Hard Silica (SiO2) Rigid network solid Insoluble in water Occurs naturally in quartz. Sand is usually small fragments of quartz. The golden-brown color is caused by iron oxide impurities. Silica gets its strength from its covalent bonding network structure. Zircon (ZrSiO4) Used as a substitute for diamonds in costume jewelry. Germanium and Tin (Ge & Sn) Germanium is recovered from the flue dust of industrial plants processing zinc ores (it occurs as an impurity in zinc)..Germanium is mainly used in the semiconductor industry Tin is easily obtained from its ore (cassiterite (SnO2)) by reduction.with carbon SnO2(s) + C(s) 1200° C Sn(l) + CO2(g) Tin is expensive and not very strong, but it is resistant to corrosion. Its main use is in tin plating and used in alloys such as bronze. Germanium - Periodic Table of Videos - YouTube Tin (version 1) - Periodic Table of Videos - YouTube  Lead Pb Lead is also easily obtained from its ore (galena (PbS)) and converted to its oxide and then reduce with carbon. 2PbS(s) + 3O2(g) ∆ 2PbO(s) + 2SO2(g) PbO(s) + C(s) Pb(s) + CO(g) Lead is durable and malleable which makes it useful in the construction industry. In addition, lead is very dense which makes it ideal as radiation shields from X-rays. Lead is also used as electrodes for rechargeable batteries. Lead - Periodic Table of Videos - YouTube Common Reactions Reaction with Halogens M +2X2 MX4 X2 = any group 17 molecule, M = Ge or Sn; Pb gives PbX2 Reactions with O M + O2 MO2  Group 15 The Nitrogen Family Group 15: The Nitrogen Family Electron configurations ns2np3. Oxidation states that range from - 3 to + 5. The metallic character of the group increases down the group. Melting and boiling points of the Group 15 elements The Element (Nitrogen): Rare in the Earth’ s crust but elemental nitrogen (N2) is the principal component of our atmosphere (76% by mass). N ≡ N triple bond strength is 944 kJ.mol-1 making it almost as inert as the noble gases. Nitrogen is used in medicines, fertilizers, explosives, and plastics. The biggest commercial use for elemental nitrogen gas is for the formation of ammonia in the Haber process. Nitrogen - Periodic Table of Videos - YouTube What Is The Haber Process | Reactions | Chemistry | FuseSchool – YouTube The chemical reaction that feeds the world - Daniel D. Dulek – YouTube Nitrogen Compounds: Ammonia (NH3) Properties: Pungent NH3 is a reasonably strong Lewis base. Toxic Gas NH3 salts decompose when heated. Condenses to clear liquid at - 33oC NH4NO3 The higher temperature reaction has explosive power and that is the reason that NH4NO3 is used as a component of dynamite. Plants need nitrogen to grow but N2 is so stable that the plants can not break the triple bond to be able to utilize the nitrogen. NH4NO3 has a high concentration of N and dissolves in water therefore it is used as a fertilizer. Rarely seen movie of Texas City, Texas Explosion - April 16, 1947 – YouTube  Hydrazine (NH2NH2) Properties: Oily Uses: Colorless Rocket Fuel Liquid Dangerous N2H4(aq) + O2(g) N2(g) + 2H 2O(l) Explosive Azide ion (N3-) Highly reactive polyatomic anion. Its most common salt is sodium azide (NaN3). Like most of the azide salts, NaN3 it is shock sensitive. NaN3 is used in airbags where it decomposes Airbags | How do they work? – YouTube to elemental sodium and nitrogen when detonated 2NaN3(s) 2Na(s) + 3N2(g) Nitrogen Oxides Properties: Dinitrogen oxide (N2O) Tasteless Unreactive Uses: Nontoxic in small Foaming agent and propellant for whipped amounts Soluble in fat cream. Nitrogen monoxide (NO) NO (which is produced from hot airplane and automobile engines) has many harmful effects: leads to acid rain, formation of smog, as well as contributes to the destruction of the ozone layer. NO is rapidly oxidized to NO2 on exposure to air. 2NO(g) + O2(g) 2NO2(g) Nitrogen Dioxide (NO2) Brown poisonous gas that contributes to the color and odor of smog. NO2 dissolves in water to form nitric acid and nitrogen oxide which is what leads to acid rain. 3NO2(g) + H2O(l) 2HNO3(aq) + NO(g) Nitric acid (HNO3) HNO3 is used in the production of fertilizers and explosives. It is both an acid and an oxidizing agent. It is made in the three-step Ostwald process. STEP 1: Oxidation of ammonia 4NH3(g) + 5O2(g) 850o ,5 , / 4NO(g) + 6H2O(g) STEP 2: Oxidation of nitrogen oxide 2NO(g) + O2(g) 2NO2(g) STEP 3: Disproportionation (single atom is both oxidized and reduced) in water; 3NO2(g) + H2O(l) 2HNO3(aq) + NO(g) The Element (Phosphorus) The radius of phosphorus is nearly 50% bigger than that of nitrogen. Thus, P is too big to approach each other close enough for their 3 orbitals to overlap and form π bonds. The availability of the 3 orbitals means that phosphorus can form as many as six bonds. Condensed phosphorus vapor is called white phosphorus and is a soft, white, poisonous. Phosphorus - Periodic Table of Videos - YouTube White phosphorus is terrifying – YouTube What are white phosphorus bombs? – YouTube White phosphorus changes to red phosphorus (amorphous network) when heated in the absence of air. Red phosphorus is much less reactive. Red phosphorus is used in the striking surfaces of matchbook because the phosphorus ignites with friction.  Phosphoric acid (H3PO4) Used primarily to produce fertilizer, food additives, and detergent. Many soft drinks owe their tart taste to the presence of a small amount of phosphoric acid.  Group 16 The Oxygen Family Group 16: The Oxygen Family Electron configurations ns2np4. Melting and boiling points of the Group 16 elements  Oxygen O2 Properties: Colorless The free element accounts for 23% of the mass of Tasteless the atmosphere. Odorless Oxygen is much more reactive than nitrogen. Condenses to a pale blue liquid The most common form of elemental oxygen is O2. The biggest consumer of oxygen in industry is the steel industry which needs about 1 t of oxygen to produce 1 t of steel. In steelmaking, oxygen is blown into molten iron to oxidize any impurities, particularly carbon. O2 is also used for welding and in medicine. https://www.youtube.com/watch?v=WuG5WTId-IY Ozone (O3) Properties: O3 is formed in the stratosphere by the effects of Blue gas solar radiation on O2 molecules. Condenses O3 is present in smog where it is produced by the at following reaction: - 112oC O(g) + O2(g) O3(g) Note the O(g) is produced by O2(g) O(g)  Sulfur S Sulfur behaves differently than oxygen due to its increased size and decreased electronegativity. O can form H bonds while S cannot. Sulfur also has weaker tendencies to form multiple bonds to one atom. Instead, it can extend its octet by using its orbitals and form as many as six bonds to separate atoms. https://www.youtube.com/watch?v=mGMR72X8V-U Sulfur S Sulfur has a striking ability to catenate, or forms chains of atoms. Oxygen’ s ability to form chains is limited. Sulfur, S8  Sulfur S Sulfur is found in many types of ores. Because the ores are so common, sulfur is usually obtained as a by-product of the extraction of several metals (most notably Cu). Sulfur has a low melting point. To extract the sulfur a process called the Frasch process is used. The Frasch process entails using super heated water to melt the solid sulfur and then uses compressed air to push the resulting slurry out. Frasch Process for Extraction of Sulphur - YouTube Sulfur S Uses: Properties: Most sulfur is used to make sulfuric acid. The other largest use of sulfur is to Yellow vulcanize rubber. Tasteless Almost Odorless Insoluble Nonmetallic Vulcanization of Rubber | 12th Std | Chemistry | Science | CBSE Board | Home Revise - YouTube The Story of Vulcanized Rubber: Goodyear's Remarkable Discovery - YouTube Selenium (Se) and Tellurium (Te) Selenium and tellurium occur in sulfide ores. They are also recovered from the refining of copper. The conductivity of selenium is increased by exposure to light and so it is used in solar cells, photoelectric devices, and photocopying machines. Selenium - Periodic Table of Videos - YouTube Tellurium - Periodic Table of Videos - YouTube Oxygen Compounds with Hydrogen The most important compound of O and H is water, H2O. Water is reactive compound and considered aggressively corrosive..H2O is an oxidizing agent 2H2O(l) + 2e-  2OH-(aq) + H2(g) E = -0.42 V at pH = 7 H2O is also a mild reducing agent 4H+(aq) + O2(g) + 4e-  2H2O(l) E = 0.82 V at pH = 7 Oxygen Compounds with Hydrogen Hydrogen Peroxide (H2O2) The presence of the second oxygen atom in H 2O2 as apposed to H2O makes H2O2 a very weak acid (pKa1 = 11.75 ). H2O2 is also a stronger oxidizing agent than water. H 2O2 is sold for industrial uses as a 30% by mass aqueous solution. A 6% H2O2 solution acts to oxidize the pigments in hair in order to bleach it A 3% H2O2 solution acts a a mild antiseptic. Oxygen Compounds with Hydrogen Except for H2O, all the other Group 16 binary compounds with hydrogen (H2E where E is a group 16 element ) are toxic gases with offensive odors. They are insidious poisons because they paralyze the olfactory nerve and soon after exposure the victim cannot smell them. Example: Hydrogen sulfide (H2S) smells like rotten eggs because egg proteins contain sulfur and eggs give off the gas when they decompose.  Sulfur Oxides and Oxoacids Sulfur forms several oxides that in atmospheric chemistry are referred to collectively as SOx (read “sox” ) Propertie Sulfur dioxide (SO2 )  s: Uses:  Gas Colorless Refrigerant, preserve dried fruit, bleach for textiles and flour, Choking producing sulfuric acid. Poisonous SO2 is the starting material for making sulfur trioxide 2SO2(g) + O2(g) 2SO3(g) The SO3 is then used to make sulfuric acid.  Sulfur Oxides and Oxoacids Sulfuric acid (H2SO4) It can be produced very cheaply. H2SO4 is the most heavily produced inorganic chemical worldwide. Properties: Colorless Corrosive Oily Uses:  liquid It is widely used in industry for the production of fertilizers, petrochemicals, and detergents.  Sulfur Oxides and Oxoacids Sulfuric acid (H2SO4) The powerful dehydrating ability of sulfuric acid can be seen when a little concentrated acid is poured on sucrose (C12H22O11) C12H22O11(s) 12C(s) + 11H2O(g) CO and CO2 generated in a side reaction cause the froth. CARBON TOWER from Table Sugar - Dehydration by Sulfuric acid - Jeffrey Vinokur, Discovery Channel - YouTube  Sulfur Halides Properties: Sulfur reacts directly with all the halogens except iodine. Dense Colorless Sulfur hexafloride (SF6)  Odorless Thermally stable Sulfur reacts spontaneously in fluorine and burns brightly to Nontoxic gas give SF6. Despite its high oxidation number (+6), it is not a good oxidizing agent. It is a good insulator in air and is used in switches on high-voltage power lines. Sulfur Halides Disulfur dichloride (S2Cl2) S2Cl2 is one of the products of the reaction of sulfur with Properties: chlorine. Yellow Uses: Liquid Vulcanization of rubber. Nauseating Smell When S2Cl2 reacts with ethene (C2H4), mustard gas is formed which has been used in chemical warfare.  Group 17 The Halogens  Group 17: The Halogens Electron configurations ns2np5. In its elemental state, all halogens atoms combine to form diatomic molecules (ex F2,I2,...). Except for F, the halogens can also lose valence electrons and their oxidation states can range from -1 to +7. Melting and boiling points of the Group 17 elements  Fluorine F2 Fluorine is the halogen with greatest abundance in the Properties Earth’ s crust. : Fluorine is the most strongly oxidizing element. Therefore, Colorless it cannot be obtained from its compounds by oxidation with Gas another element. Most of the F produced by industry is used to make the volatile solid UF6 used for processing nuclear fuel. The next biggest use of F is the production of SF6 for electrical equipment. https://www.youtube.com/watch?v=vtWp45Eewtw  Chlorine Cl2 Chlorine is more soluble in water than fluorine. As a result, even though there is more F present in Properties the Earth’ s crust the oceans are salty with chlorides : rather than fluorides. Cl is one of the most heavily manufactured Pale yellow chemicals. Gas It is obtained from electrolysis of molten rock salt (NaCl) or brine. It is a strong oxidizing agent. https://www.youtube.com/watch?v=BXCfBl4rmh0  Chlorine Cl2 Uses:  In several industrial processes, including the manufacture of plastics, solvents, and pesticides. It is also used as bleach in the paper and textile industries and as a disinfectant in water treatment plants. In addition, Cl is used to produce Br. Bromine Br2 :Uses  Br is used widely in synthetic organic chemistry because of the ease at which it can be added to and Properties removed from organic chemicals that are being used : to carry out complicated syntheses. Corrosive Organic bromides are incorporated into textiles as Red-Brow fire retardants and are used as pesticides. Inorganic n Liquid bromides, particularly silver bromide, are used in photographic emulsions. https://www.youtube.com/watch?v=Slt3_5upuSs Iodine I2 When iodine dissolves in organic solvents it produces solutions having a variety of colors. These colors arise from the different interaction between the I2 molecules and the solvent. Iodine is an essential trace element for living systems; a deficiency in humans leads to a swelling of the thyroid gland in the neck. Iodides are added to table salt (iodized salt) to prevent this deficiency. https://www.youtube.com/watch?v=JUBsJLRSM64 Compounds of the Halogens The halogens form compounds among themselves. These inter halogens have the formulas XX’ , XX’ 3 , XX’ 5, and XX’ 7 (X heavier halogen) These compounds are prepared by direct reaction of the two halogens, the product formed being determined by the proportions of the reactants used. Example: Cl2(g) + 3F2(g) 2ClF3(g) Cl2(g) + 5F2(g) 2ClF5(g)  Group 18 The Nobel gases Group 18 The Nobel gases Electron configurations ns2np6. Their closed shell electron configuration makes them have a very low reactivity. Boiling points and density of the Group 18 elements  The Elements All the noble gases occur in the atmosphere as monatomic gases. Together they make up 1% (by mass) of the atmosphere. Argon is the third most abundant gas in the atmosphere after nitrogen and oxygen. All the noble gases except He and Rn are obtained by the fractional distillation of liquid air. Noble Gases - The Gases In Group 18 | Properties of Matter | Chemistry | FuseSchool - YouTube Helium He Helium is the second most abundant element in the universe after hydrogen. However, it is rare on earth because it is so light that it can reach the high speeds needed to escape from the atmosphere. Helium is found as a component of natural gases trapped under rock formations where it has collected as a result of the emission of α particles by radioactive elements. Helium gas is twice as dense as hydrogen under the same conditions. Its density is still very low, and it is nonflammable therefore it is used to provide buoyancy in blimps. https://www.youtube.com/watch?v=M6xZZiaLOV4  The Elements Neon glows orange-red when an electrical current is passed through it and is used for advertising signs and displays. Argon is used to provide an inert atmosphere for welding to prevent oxidation. Argon is also used to fill some types of light bulbs, where it conducts heat away from the filament. Krypton gives an intense white light when an electrical current is passed through it, and it is used in airports for runway lights. Xeon is used in halogen lamps, for automobile headlights, and in high-speed photographic flash tubes. Radon is a radioactive gas that seeps out of the ground and its presence can lead to dangerously high levels of radiation. The Elements Neon: https://www.youtube.com/watch?v=wzv0pb7mzaw Argon: https://www.youtube.com/watch?v=N0Gw6-xMLlo Krypton: https://www.youtube.com/watch?v=il4OOY7Zseg Xeon: https://www.youtube.com/watch?v=Ejoct_6pQ74 Radon: https://www.youtube.com/watch?v=mTuC_LrEfbU  Compounds of the Nobel Gases The ionization energies of the noble gases are very high but decrease down the group. No compounds of helium, neon, or argon exist except under very special conditions. Krypton forms only one known stable neutral molecule KrF2. Xenon’ s ionization energy is low enough for electrons to be lost to very electronegative elements. Xe forms several compounds with fluorine and oxygen and compounds with Xe-N and Xe-C bonds have been reported. Xenon fluorides are used as powerful fluorinating agents (reagents for attaching fluorine atoms to other substances).  Ionic Bonds / Ionic Compounds Covalent Bonds / Covalent Compounds Why do atoms bond?  Each atom wants a full outermost energy level. Gain, lose, and share valence electrons to achieve the duet or octet. Gives each atom an electron configuration similar to that of a noble gas  ex. Group 18: He, Ne, Ar Chemical Bonds Attractive force that holds atoms or ions together. :types 3  ionic, covalent, metallic Determines the structure of compound. Structure affects properties:.melting/boiling points, conductivity etc -  Chemical Structure/Molecular Models Arrangement of bonded atoms or ions. Bond length: the average distance between the nuclei of two bonded atoms. Bond angles: the angle formed by two bonds to the same atom. Ball and stick Atoms are represented by balls. Bonds are represented by sticks. * Good for “seeing” angles. Structural Chemical symbols represents atoms. Lines are used to represent bonds. * Good for “seeing” angles.  Space filling Colored circles represent atoms, and the space they take up. No bonds, no bond angles. Electron Dot/Lewis Structure Chemical symbol represent atom. Dots represent valence electrons. 2 center dots represent a bond. No bond angles, no bond length. Ionic Bonds / Ionic Compounds Definition Bond formed by the attraction between oppositely charged ions cation: positive: lost e-’ s anion: negative: gained e-’ s EX: Na (sodium) needs to lose one electron to become stable, Cl (chlorine) needs to gain one electron to become stable. Na becomes positive, Cl becomes negative and they bond due to their opposite charges. Electrons are transferred from one atom to another. Negative ions attract more positive ions, and soon a network is formed.  Networks Repeating pattern of multiple ions, ex. NaCl Every Na ion is next to 6 Cl ions Strong attraction between ions creates a rigid framework, or lattice structure (crystals). ex, cubic, hexagonal, tetragonal. Properties of Ionic Compounds Structure affects properties Strong attractions between ions: strong bonds. High melting/boiling points. Shatter when struck (think of it as one unit). Conductivity: Solid: ions are so close together, fixed positions, (can’ t move)  NO conductivity. Liquid: ions are freely moving due to a broken lattice structure  Good conductivity. Writing the Formula When ionic compounds form, they balance out the charges of the ions. The formula must represent this balance. 2 Chloride ions (-1) will balance out the charge of a Magnesium ion (+2). We write this formula out: MgCl2 The subscript 2 tells us that we have 2 Chlorine atoms. If no subscript is written that means, there is only one atom. Naming Ionic Compounds Always name the cation (positively charged ion) first, the anion (negatively.1 charged ion) second. Use the name of the cation without any alterations..2 Use the root and “-ide” to the end for anions..3 For the transition metals, use brackets and Roman Numerals after the name.4 to denote its charge. Examples LiI Lithium Iodide FeCl2 Iron (II) Chloride BeBr2 Beryllium Bromide Al2S3 Aluminum Sulfide FeCl3 Iron (III) Chloride Polyatomic Ions Ions formed of multiple atoms usually covalently bonded together and should be treated as one entity. Nitrate: NO3- Pb2+ + 2NO3- → Pb(NO3)2 The brackets tell us that everything inside is doubled. Common Polyatomic Ions NH4+ Ammonium NO2- Nitrite CO32- Carbonate NO3- Nitrate PO43- Phosphate SO32- Sulfite ClO2- Chlorite SO42- Sulfate ClO3- Chlorate MnO4- Permanganate ClO4- Perchlorate OH- Hydroxide HCO3- Hydrogen carbonate (Bicarbonate) Covalent Bonds / Covalent Compounds Definition  Chemical bond in which two atoms share a pair of valence electrons. Can be a single, double, or triple bond. Single, 2e-’ s; double, 4e-’ s; triple, 6e-’ s always formed between nonmetals -.mostly low melting/boiling points - 2 types of bonds  Polar Non polar Non Polar  Bonded atoms that share e-’ s equally - Same atoms bonded - ex. Cl – Cl: Cl2 Polar  bonded atoms that do not share e-’ s equally - different atoms bonded - H ex. H – N – H: NH3 Properties of Covalent Compounds Covalent compounds generally have much lower melting and boiling points than ionic compounds. Covalent compounds are soft (compared to ionic compounds). On the other hand, covalent compounds have these molecules which can very easily move around each other, because there are no bonds between them. As a result, covalent compounds are frequently flexible rather than hard. Properties of Covalent Compounds Covalent compounds tend to be more flammable than ionic compounds. Covalent compounds don't conduct electricity in water. Covalent compounds aren't usually very soluble in water. How is the ratio of atoms calculated? To calculate the ratio of atoms in a stable covalent compound: 1. Work out how many electrons are needed by each non-metal element to complete its outer electron shell. 2. Work out the ratio of atoms that will provide enough shared electrons to fill all the outer shells. For example, how H N Elements many nitrogen and 1 2.5 Electron configuration hydrogen atoms bond 1 3 Electrons needed together in an 3 1 Ration atoms ammonia molecule? How do carbon and hydrogen atoms form covalent bonds in a molecule of methane? H C Elements 1 2.4 Electron configuration 1 4 Electrons needed 4 1 Ration atoms H CH4 or H C H H How do carbon and oxygen atoms form covalent bonds in a molecule of carbon dioxide? O C Elements 2.6 2.4 Electron configuration 2 4 Electrons needed 2 1 Ration atoms A double bond is when two pairs of electrons are CO2 or O C O shared. In carbon dioxide there are two double bonds one between each oxygen atom and the carbon atom. Naming Covalent Compounds Name the first non-metal If there is only one atom present, DO NOT use a prefix. If there is more than one atom present, then you must use the correct prefix. Name the second non-metal  Using the correct prefix and an -ide ending. Prefixes-stand for atoms present 1 -- mono¸ 6 -- hexa¸ 2 -- di¸ 7 -- hepta ¸ 3 -- tri¸ 8 -- octa ¸ 4 -- tetra¸ 9 -- nano¸ 5 -- penta ¸ 10 -- deca¸ Examples: CO2 Carbon Dioxide N3 Trinitrogen N2O5 Dinitrogen Pentoxide CCl4 Carbon Tetrachloride Cl2O Dichlorine Monoxide CO Carbon Monoxide PCl3 Phosphorus Trichloride CI4 Carbon Tetriodide Metallic Bonds Definition  A bond formed by the attraction between positively charged metal ion (cation) and the shared electrons that surround it (sea of electrons) ex. Cu Properties  Conductivity: Good: electrons can move freely. Malleable: lattice structure is flexible. Classification of Bonds You can determine the type of bond between two atoms by calculating the difference in electronegativity values between the elements Electronegativity Difference Type of Bond 0.4 – 0 Nonpolar Covalent 1.9 – 0.5 Polar Covalent 4.0 – 2.0 Ionic Chemical Intermolecular Forces Intermolecular forces Intramolecular forces hold atoms together in a molecule. Covalent bonds would be an example. Intermolecular forces are attractive forces between molecules. Intramolecular = strong Intermolecular = weak They do control physical properties such as boiling and melting points, vapor pressure, and viscosity. Types of Intermolecular Forces These intermolecular forces as a group are referred to as van der Waals forces Van der Waals Forces  Dipole-dipole interactions.(example Hydrogen bonding). Ion-Dipole interactions. London dispersion forces. Dipole-dipole interactions Molecules that have permanent dipoles  are attracted to each other. The positive end of one is attracted to the negative end of the other and vice versa. These forces are only important when the molecules are close to each other. Ion-Dipole interactions Attractive forces between an ion and a polar molecule The larger the charge the stronger the force  A molecular picture showing the ion-dipole Interaction that helps a solid ionic crystal dissolve in water.  London dispersion forces The London dispersion force is the weakest intermolecular force. While the electrons in the 1s orbital of helium would repel each other (and, therefore, tend to stay far away from each other), it does happen that they occasionally wind up on the Intermolecular same side of the atom. At that instant, then, the helium atom is polar, with an excess of electrons on the left side and a shortage on the right side.  London dispersion forces Another helium nearby, then, would have a dipole induced in it, as the electrons on the left side of helium atom 2 repel the electrons in the cloud on helium atom 1. London dispersion forces, or dispersion forces, are attractions between an instantaneous dipole and an induced dipole. These forces are present in all molecules, whether they are polar or nonpolar. The tendency of an electron cloud to distort in this way is called polarizability. The larger the molecule the greater it’ s Dispersion Forces are. Hydrogen bonding The dipole-dipole interactions experienced when H is bonded to N, O, or F are unusually strong. We call these interactions hydrogen bonds Hydrogen bonding arises in part from the high electronegativity of nitrogen, oxygen, and fluorine. Also, when hydrogen is bonded to one of those very electronegative elements, the hydrogen nucleus is exposed. Water in the liquid and solid states exists as groups in which the water molecules are linked together by hydrogen bonds. Summarizing Intermolecular Forces