Foundation Programme Chemistry 2022 PDF

Foundation programme Chemistry 2022 Petr Táborský 1 TODAY'S TOPICS Ionic bond. Formation of ions. Ionic charges from group number. Metallic bond. Names and formulas of ionic compounds. 2 ...

Foundation programme Chemistry 2022 Petr Táborský 1 TODAY'S TOPICS Ionic bond. Formation of ions. Ionic charges from group number. Metallic bond. Names and formulas of ionic compounds. 2  Ionic bond ionic bond is the extreme form of the polar covalent bond. the larger the difference in electronegativity between the two atoms involved in the bond, the more ionic (polar) the bond is. the shared electrons are pulled into the vicinity of the more electronegative atom, so that 2 ions are created. The cation and anion with opposite charges are attracted by electrostatic forces. Electronic Electronic Electronic Electronic configuration configuration configuration configuration 2,8,8,1 2,8,7 2,8,8 2,8,8 xx xx x x x Cl- x Na x Cl- Na+ x x x 3 xx xx  Formation of ionic bond By losing one electron, sodium atom achieves the electronic configuration [2,8]. By gaining one electron, chloride atom achieves the electronic configuration [2,8]. [2,8] is the electronic configuration of neon; it is a ‘noble- gas configuration’. 4  Polarity of bonds The polarity of the bond can be approximately derived from the difference of electronegativities of both bound elements. The generally used rule says -if the difference < 0.4 non- polar bond -if the difference is 0.4–2.0 polar bond - if the difference is higher 2.0 ionic bond There is no sharp dividing line between ionic and covalent bonds. 5  Ionic bond Ionic bond is typically formed between atoms of metals and nonmetals 6  Example: assess the bond polarity in the following compounds: Compound F2 HF LiF Electronegativity difference Type of bond Electronegativities: H – 2,2 Li – 1,0 F – 4,0 7  Ionic structure of magnesium oxide Mg + O MgO EN 1,3 EN 3,4 Electronic Electronic Electronic Electronic configuration configuration configuration configuration 2,8, 2 2,6 2,8,8 2,8,8 xx xx x x x x Mg O Mg2+ O2- x x x x xx By losing two electrons, each magnesium atom achieves the electronic configuration [2,8]. By gaining two electrons, each oxygen atom achieves the electronic configuration [2,8]. [2,8] is the electronic 8 configuration of neon; it is a ‘noble-gas configuration’.  Explain the formation of ionic structure of EN … CaCl2 EN … Cl Ca Cl 9 Which of the following pairs of elements are likely to form an ionic compound? a) Helium and oxygen b) Chlorine and bromine c) Magnesium and chlorine d) Nitrogen and iodine e) Potassium and sulfur 10 Write the symbols for the ions and correct formula for the ionic compound formed by each of the following: a) Magnesium and oxygen b) Potassium and sulfur c) Silicon and oxygen d) Phosphorus and bromine e) Aluminium and chlorine 11  Ionic compounds composed of metal and nonmetal one or more electrons are transfered from metal to nonmetal crystalline substances in solid state. the ions are bound by strong electrostatic forces in the crystals; high melting and boiling points. most of ionic compounds are good soluble in water and their solutions and melts conduct electric current (electrolytes). in the solid state there are no free molecules, they form ion lattices. the chemical formula, e.g. NaCl, is only the simplest record of the stoichiometric ratio of ions in the crystal.  Covalent compounds Two or more nonmetals that share one or more valence electrons much less polar than the ionic compounds. They include mainly molecular substances, where the chemical formula really expresses the smallest particle of the compound (CO2, CH3OH). They occur as gases, liquids, and solids with relatively low melting and boiling temperature. 13 Label each of these bonds as covalent, polar covalent, or ionic: a) K—Cl b) Br—Cl a) Cl—Cl 14 Names and Formulas of Ionic Compounds 15  Names and charges of Cations The name of a cation consists of the name of the element and the word ion (or cation). Fix the charges of the most common cations Charge 1+ Charge 2+ Charge 3+ Li+ lithium Mg2+ magnesium Al3+ aluminium Na+ sodim Ca2+ calcium K+ potassium Sr2+ strontium NH4+ ammoniu Ba2+ baryum m Ag+ silver Zn2+ zinc Cd2+ cadmium 16 Why cations of alkaline metals have always charge 1+? Why cations of alkaline earth metals have always charge 2+? 17 Monoatomic ions and their nearest noble gases Metals lose Nonmetals gain valence electron valence electrons Noble 1A 2A 3A 5A 6A 7A Noble gases gases He Li+ Ne Na+ Mg2+ Al3+ N3- O2– Cl– Ne Ar K+ Ca2+ P3- S2– F– Ar Kr Rb+ Sr2+ Br– Kr Xe Cs+ Ba2+ I– Xe 18 State the number of electrons that must be lost by atoms of each of the following to achive a noble gas configuration: Na Mg Al Sr K 19  Metals with variable charges We cannot determine the charge of transition element from their group number They typically form two or more positive ions The transition elements lose electrons , but they are lost from highest energy level and sometimes from a lower energy level as well. This is also true for metals of representative elements in Group 4A and 5A e.g.Pb, Sn, Bi For metals that form two or more ions, a naming system is used to identify the particular cation To do this, a Roman numeral inside parenthesis denotes the oxidation number E.g.Fe3+is Fe(III), Fe2+ is Fe(II) 20  Names and charges of cations Charge 1+ or 2+ Cu+ Copper (I) Cu2+ Copper (II) Hg22+ Mercury (I) Hg2+ Mercury (II) Charge 1+ or 3+ Au+ Gold (I) Au3+ Gold (II) Charge 2+ or 3+ Fe2+ Iron (II) Fe2+ Iron (III) Co2+ Cobalt (II) Co3+ Cobalt (III) Cr2+ Chromium (II) Cr3+ Chromium (III) Mn2+ Manganese (II) Mn3+ Manganese (III) Ni2+ Nickel (II) Ni3+ Nickel (III) Charge 2+ or 4+ Sn2+ Tin (II) Sn4+ Tin (IV) Pb2+ Lead (II) Pb4+ Lead (IV) There is an older rule still used, by which the names of elements are transformed into adjectives by combining the Latin stem of element’s name with the suffixes –ous for the lesser charge and –ic for the greater charge. Fe2+ iron(II) ion Cu+ copper(I) ion ferrous ion cuprous ion Fe3+ iron(III) ion Cu2+ copper(II) ion ferric ion cupric ion 22  Names and charges of Anions The name of a monatomic anion consists of adjective formed from the stem of the name of the element modified with the suffix –ide and the noun ion Charge 1- Charge 2- Cl– chloride ion (-I) O2– oxide ion (-II) F– fluoride ion (-I) S2– sulfide ion (-II) I– iodide ion (-I) Charge 3- Br– bromide ion (-I) N3- Nitride ion (III) H– hydride ion (-I) P3- Phosphide (III) 23 Why anions of halogenes (F,Cl, Br,I) have charge -1 ? 24 Monoatomic ions and their nearest noble gases Metals lose Nonmetals gain valence electron valence electrons Noble 1A 2A 3A 5A 6A 7A Noble gases gases He Li+ Ne Na+ Mg2+ Al3+ N3- O2– Cl– Ne Ar K+ Ca2+ P3- S2– F– Ar Kr Rb+ Sr2+ Br– Kr Xe Cs+ Ba2+ I– Xe 25 State the number of electrons that must be gained by atoms of each of the following to achive a noble gas configuration: F O P As S 26 State the number of electrons that must be gained or lost when the following elements form ions: Be Sb Cl Ba I 27  Chemical formulas of ionic compounds The chemical formulas of a compound represents the symbols and subscripts in the lowest whole number ratio of the atoms or ions In the formula of an ionic compound the sum of ionic charges in the formula is always zero. Thus the total amount of positive charge is equal to the total amount of negative charge. 28 Consider a compound containing Mg2+ and Cl- Solution: To balance the ionic charge of the Mg2+ we need to place 2 Cl- ions in the formula: This gives the formula MgCl2 with the overall charge zero 29 Write the symbols for ions, and the correct formula for the ionic compound formed when sodium reacts with sulfur Solution: Sodium, which is a metal of group IA, forms Na+ cation; Sulfur, which is nonmetal in group 6A, forms S2- The charge of 2- is balanced by two Na+ ions. The ionic formula of the compoud is Na2S 30 Write the name for the ionic compound CaF2 Solution: Identify the cation and anion The cation Ca is from the group 2A, the anion F is from group 7A Name the cation: calcium Name the anion by using the name of the first syllable of the element name followed by –ide: fluoride Write the name for the cation first and the name for the anion second: calcium fluoride 31  Determination of cation charge transition elements When you name an ionic compound , you need to determine if the metal is a representative element or a transition element If it is transition element , except for cadminum, zinc and silver, you will need to use its ionic charge as a Roman numeral as a part of its name The calculation of ionic charge depends on the negative charge of the anions in the formula 32 What is the name of ionic compound CuCl2? Solution: We know that charge of Cl ion is -1, two chlorides have -2, the copper ion must have charge 2+ (Cu2+) The name of the compound is copper(II) chloride 33  Write the formula of iron (III) chloride: Write the formula of iron (III) sulfide: 34  Give the structures of the following compounds Name Formula Magnesium chloride Potassium sulfide Baryum nitride Iron (III) oxide Calcium sulfide Ammonium sulfide Calcium phosphide Iron(II)bromide Iron(III) iodide 35 Repetition 36 TODAY'S TOPICS Weak intermolecular bonding. Naming covalent compounds, writing formulas. 37  Intermolecular Forces the attractive forces between molecules (or between the parts of a macromolecule), which influence the consistence of substances. The energy of these interactions is much lower than the energy of the covalent bonds and the ionic interactions in the crystal lattices they are defined as the weak non-bonding interactions or non-covalent bonds. 38  Intermolecular Forces The character and strength of the intermolecular forces determines e.g. the state of the given substance and affects its solubility. Their biological importance: - e.g.n maintaining the secondary, tertiary or quaternary structure of biopolymers, - stability of the supermolecular structures such as biomembranes, - in the specific biological interactions (the binding of a substrate to an enzyme, an antibody to antigen, a hormone to protein receptor) etc. 39 Classification of intermolecular forces Hydrogen Bonds Van der Waals Forces Electrostatic interactions Dispersion forces Hydrophobic interactions 40 Energy of intermolecular forces – comparision with ionic and molecular bonds Type of bond Bond strength ( kJ mol–1) hydrogen bonding 20–50 dipole–dipole force 5–20 van der Waals’ forces 1–20 Ionic bond 700-4000 Covalent bond 200-1000 41  Hydrogen Bonds the strongest intermolecular forces they are found in substances, in which the hydrogen atom is bonded to a strongly electronegative atom – nitrogen, oxygen, or fluorine. O-H, F-H, N-H The electronegative atom of oxygen strongly attracts the shared electron pair Hydrogen bond the small hydrogen atom H has little electron density - - around it. + Under these circumstances, O H O R the hydrogen atom H carries a partial positive charge and is able to form a weak bond Electrons atracted Electrons atracted with a free electron pair of the electronegative oxygen of the other molecule 42  Hydrogen Bonds The bond energy is greater than in the case of other non-covalent interactions, it reaches up to several tens of kJ mol−1. The hydrogen bonds can be intermolecular and intramolecular. R HC C O C O H N HN CH R 43 intermolecular intramolecular Hydrogen bonds between the molecules of water Hydrogen bonds have special importance for the properties of water. Because oxygen in H2O has two lone pairs of electrons and two covalently bonded hydrogen atoms, each water molecule is able to form hydrogen bonds with up to four other molecules at the same time. Thus aggregates (clusters) with a different number of molecules and with circular or spatial arrangement are formed. They have tetrahedral arrangement in ice – each molecule binds four others and six-member rings are formed. The crystal lattice is relatively loose (lower 44 density of ice, increase of volume). Molecules of water have tetrahedral arrangement in ice – each molecule binds four others and six-member rings are formed. The crystal lattice is relatively loose (lower density of ice, increase of volume). Expansion of water upon freezing 45 The hydrogen bonds can be intermolecular and intramolecular Intramolecular bridges between peptide chains 46 Hydrogen bonds Hydrogen bonds in DNA in - helix of proteins double helix 47  Van der Waals Forces Electrostatic interactions Dispersion forces Hydrophobic interactions 48  Electrostatic interactions The principle are attractive coulombic forces between the positive and negative electrical charges. Electrical charge can be either whole, in ions, or partial, in dipoles. Ion-ion interactions + - Ion-dipole + - dipole-dipole interactions + - 49  Ion – ion interaction ion - + ion formation of ion pairs (salt bridges) in solutions In proteins during interaction of side amino acid chains They cannot be confused with ionic bonds in the crystal lattice of ionic compounds  Dipol ion, dipol dipol interactions Interactions between polar molecules The partially positive ion + interacts with partially negative - ion solvation of ions in polar solvents stabilization of the tertiary structure of proteins. solubility of polar substances in polar solvents 51 Other examples of dipol- dipol or ion- dipol interaction Hydration of Na+ Hydration of Cl- 52  Dispersion forces Dispersion forces occur in non-polar substances (carbohydrates, gaseous elements including noble gases). they are the result of momentary shifts in the symmetry of electron cloud in the molecule. In a large collection of molecules, at any given moment collisions are taking place, with resulting polarization of the molecules. 53  Dispersion forces As soon as a slight positive charge is produced at one end of the molecule, it induces a slight negative charge in one end of the molecule next to it – induced dipole. Dispersion forces are attraction between fluctuating dipoles in atoms and molecules that are very close together. 54  Dispersion forces In nonpolar substances (hydrocarbons, O2, N2, H2, inert gases), the dispersion forces are the only intermolecular interactions. They are responsible for the condensation of gases at low temperature. The bond energy induced by the dispersion forces is 2 or 3 orders lower than the energy of the covalent or ion bonds. 55 Dispersion forces increase with increasing number of electrons (and protons) in the molecule increasing the number of contact points between the molecules – contact points are places where the molecules come close together. 56 Boiling points of the noble gases The boiling points of the noble gases increase as the number of electrons increases. This is because the van der Waals’ forces between the atoms are increased with an increasing number of electrons. 57  Compare the boiling points of pentane and 2,2-dimethylpropane. These compounds have equal numbers of electrons in their molecules boiling point 36 °C boiling point 10 °C The molecules in pentane can line up beside each other so there are a large number of contact points. The van der Waals’ forces are higher, so the boiling point is higher. The molecules of 2,2-dimethylpropane are more compact. The surface area available for coming into contact with neighbouring molecules is smaller. The van der Waals’ forces are relatively lower, so the 58 boiling point is lower.  Hydrophobic interactions the bonds formed between the hydrophobic (non-polar) molecules of substances in an aqueous environment. When you add some drops of oil to water, the drops combine to form a larger drop. Water molecules are attracted to each other and are cohesive because they are polar molecules. Oil molecules are non- polar and thus have no charged regions on them. This means that they The attractiveness of the water are neither repelled nor molecules then has the effect of attracted to each other. squeezing the oil drops together to form a larger drop. 59 Hydrophobic oil immiscible in water 60 Complete and explain the table: Substance Melting point Type of bonds (oC) between particles MgF2 1248 Ionic bond NaCl 801 ……………. H2O …….. ………… NH3 -78 …………… HBr -89 ……………. HCl -115 …………… Cl2 -101 ……………. F2 -220 …………… 61 Indicate the major type of molecular (non covalent) interaction expected for each of the following: HF Br2 PCl3 Why is the melting point of H2S lower than that of H2O ? 62 Indicate the major type of attractive force 1 – ionic bonds 2 – dipole-dipole interaction 3 – hydrogen bond 4 – dispersion forces a) NH3 b) HI c) I2 d) Cs2O e) N2 63 Nomenclature - Polyatomic ions 64  Nomenclature - Polyatomic ions An ionic compound may also contain a polyatomic ion as one of its cations or anions. A polyatomic ion is a group of covalently bonded atoms that has an overall ionic charge Most polyatomic ions consist of nonmetal such as phosphorus, sulfur, carbon or nitrogen covalently bonded to oxygen atom 65  Names of polyatomic ions Names of polyatomic ions end in -ate, such as nitrate, sulfate etc. When a related ion has one less oxygen atom, the ending - ite is used such as nitrite, sulfite etc. The hydroxide ion (OH-) and cyanide ion (CN-) are exceptions to this naming pattern 66 Names and formulas of some common polyatomic ions Nonmetal Formula of ion Name of ion Hydrogen OH- hydroxide Nitrogen NH4+ ammonium NO3- nitrate NO2- nitrite chlorine ClO4- perchlorate ClO3- chlorate ClO2- chlorite ClO- hypochlorite carbon CO32- carbonate HCO3- Hydrogen carbonate CN- cyanide sulfur SO42- sulfate HSO4- hydrogen sulfate 67 Names and formulas of some common polyatomic ions cont. Nonmetal Formula of ion Name of ion sulfur SO32- sulfite HSO3- hydrogen sulfite Phosphorus PO43- phosphate HPO42- Hydrogen phosphate H2PO4- Dihydrogen phosphate PO33- phosphite 68  Writing formulas for compounds containing polyatomic ions Polyatomic ions must be associated with ions of opposite charge E.g. The compound sodium chlorite is composed of Na+ cations and ClO2- anions To write the correct formulas we follow the same rules of charge balance that we used for writing the formulas for simple ionic compounds The net charge equals zero. 69 Consider the formula for compound containing potassium and nitrate ions. The ions are K+ and NO3- Because one ion of each balances the charge, the formula is: KNO3 70 Consider the formula for compound containing calcium and nitrate ions. The ions are Ca2+ and NO3- To balance the positive charge 2+ on the calcium ion, two nitrate anions are needed. In the formula of the compoud parenthesis are placed around the nitrate ion, and the subscript 2 is written outside the right parentesis Ca(NO3)2 71  Naming of ionic compouds containing polyatomic ions We first write the positive ion, ussually a metal, and then we write the name for the polyatomic ion. Na2SO4 Na2SO3 Sodium sulfate Sodium sulfite It is important to recognize the polyatomic ion correctly 72  Give the formulas of the following compounds Name Formula Calcium phosphate Potassium carbonate Baryum nitrite Silver nitrate Calcium sulfite Ammonium sulfite Aluminium sulfate Iron(II)sulfate Iron(III) sulfide Baryum hydroxide 73  Give the names of the following compounds formula Name Ca(NO2)2 FePO4 Mg(HCO3)2 MgSO3 Al(NO2)2 Ca3(PO4)2 NH4NO3 KH2PO4 KClO3 Al(OH)3 74  Names of mineral acids Acids are hydrogen-containing covalent compounds that liberate hydronium (hydrated ions H+) when dissolved in water. Two types of acids can be distinguished: Acids derived from binary hydrogen compounds Oxoacids 75 Acids derived from binary hydrogen compounds The chemical bond that exists between two nonmetals is predominantly covalent. Binary compounds of hydrogen and certain nometals (group 16 and 17, sulfur and halogens) are volatile, mostly gases However, when dissolved in water these compounds have acid properties. Then they are given acid name composed of the stem of the nonmetal with the prefix hydro–, the suffix –ic, and the noun acid. Formula Binary compound Acid anion H2S hydrogen sulfide hydrosulfuric acid sulfide HF hydrogen fluoride hydrofluoric acid fluoride HCl hydrogen chloride hydrochloric acid chloride 76 HI hydrogen iodide hydroiodic acid iodide  Names of common acids and their anions Formula Name Anion H3BO3 boric acid borate The most H2CO3 carbonic acid carbonate common form of an oxygen H4SiO4 silicic acid silicate containing acid H2Cr2O7 dichromic acid dichromate has name that HNO2 nitrous acid nitrite ends with HNO3 nitric acid nitrate –ic acid H2SO3 sulfurous acid sulfite H2SO4 sulfuric acid sulfate An acid that HClO Hypochlorous acid hypochlorite contains one HClO2 Chlorous acid chlorite less oxygen HClO3 Chloric acid chlorate atom is named HClO4 (Hy)perchloric acid hyperchlorate as an – ous acid H3PO4 phosphoric acid phosphate H3PO3 phosphorous acid phosphite 77 TODAY'S TOPICS Stoichiometric calculations. Chemical equations, balancing equations, mass relations in chemical reactions. Nomenclature of covalent compounds 78  Calculations from chemical formula The chemical formula specifies: Atoms in 1 molecule Moles of each element in 1 mol of a compound Chemical formula of ethanol is C2H5OH One molecule of ethanol contains 2 atoms of carbon 6 atoms of hydrogen one atom of oxygen One mol of ethanol contains 2 mols of carbon 6 mols of hydrogen one mol of oxygen 79 Problem: How many moles of hydrogen is contained in 2.5 moles of ethanol? Solution: 1 mol of C2H5OH contains 6 moles of H 2.5 moles ? (n) As the proportion of hydrogen in ethanol is always the same, it is valid: Use Cross-multiplication 2.5/ 1 = n /6 n = 2.5/ 1 x 6 = 15 moles 80 2.5 moles of ethanol contain 15 moles of hydrogen Problem: Chemical formula of ethanol is C2H5OH. What is the relative molecular mass of ethanol, and what is the mass of 1 mol of ethanol? Ar (H) = 1.01, Ar (C) = 12.01, Ar (O) = 16.00 Solution: One molecule of ethanol contains 2 atoms of carbon 6 atoms of hydrogen one atom of oxygen (2 x 12.01) g of C (6 x 1.01 ) g of H (1 x 16.00 ) g of O The relative molecular mass Mr = (24.02 + 6.06 + 16.00) = 46.08 M (molar mass) is numerically equal to the relative molecular Mr mass, but it is expressed in grams. 81 The mass of one mol of ethanol 46.08 g Problem: What is the mass fraction (% composition) of elements in molecule of ethanol? Solution: Mass fraction w of an element is calculated as w = relative mass of the element (X)/ Mr wc = 24.02/ 46.08 = 0.5213 (= 52.13 %) wH = 6.06 / 46.08 = 0.1315 (=13.15 %) wO = 16.00 / 46.08 = 0.3472 (=34.72 %) ------------- ------------------ 46.08 100 % The percentage composition of ethanol: 52.13 % of C, 13.15 % of H, and 34.72 % of O. 82 Problem: What is the mass fraction of water in sodium sulfite heptahydrate? Solution: Formula of the crystalline salt is: Na2SO3·7 H2O There are seven molecules of water (Mr (H2O) = 18) per one formula unit. Therefore, one mol of the compound contains 18 × 7 = 126 g of H2O The relative molecular and molar masses of the salt are: Mr = 252.2 M = 252.2 g mol−1 Mass fraction of water w(H2O) = 126 (g mol−1) / 252.2 (g mol−1) = 0.4996 Water incorporated in a crystal lattice of the salt represents 50.0 % (rounded to 3 digits) of its mass. 83 Problem: Which of the following is equivalent to the higher number of moles: 1 kg of hydrogen peroxide or 1 kg of water? Solution Mr (H2O2) = 34.0 M (H2O2) = 34 g mol−1 Mr (H2O) = 18.0 M (H2O) = 18 g mol−1 The number of moles of the substance is calculated as n = m/M. Substituting the values, we get: n (H2O) = 1000 (g) / 18 (g mol−1) = 55.6 mol n (H2O2) = 1000 (g) / 34 (g mol−1) = 29.4 mol Substance amount of water is higher than that of hydrogen peroxide. 84 Types of chemical reactions Chemical reactions can be divided into several classes each having similar characteristics. 85 Composition (combination) reaction A chemical reaction in which a single substance is produced from multiple reactants. 2H2(g) + O2(g) → 2H2O(ℓ) Decomposition reaction A chemical reaction in which a single substance becomes more than one substance. 2NaHCO3(s) → Na2CO3(s) + CO2(g) + H2O(ℓ) 86 Combustion reaction A chemical reaction in which a reactant combines with oxygen to produce oxides of all other elements as products CH4(g) + 2O2 (g) → CO2(g) + 2H2O(g) Single replacement reaction Elements in a compound are replaced by other elements Zn(s) + 2HCl(aq) →H2(g) + ZnCl2(aq) 87 Double replacement reaction The positive ions in the reacting compounds switch places BaCl2(aq) + Na2SO4 (aq) → BaSO4(aq) + NaCl(aq) 88 Identify each equation as a composition reaction, a decomposition reaction, or neither. Fe2O3 + 3SO3 → Fe2(SO4)3 NaCl + AgNO3 → AgCl + NaNO3 (NH4)2Cr2O7 → Cr2O3 + 4H2O + N2 89 Which is a composition reaction and which is not? NaCl + AgNO3 → AgCl + NaNO3 CaO + CO2 → CaCO3 H2 + Cl2 → 2HCl 2HBr + Cl2 → 2HCl + Br2 90  Balancing chemical equation When chemicals react, atoms cannot be either created or destroyed. So there must be the same number of each type of atom on the reactants side of a chemical equation as there are on the products side. A symbol equation is a shorthand way of describing a chemical reaction. It shows the number and type of the atoms in the reactants and the number and type of atoms in the products. If these are the same, we say the equation is balanced. 91 Problem: Write a balanced equation for the reaction of iron(III) oxide with carbon monoxide to form iron and carbon dioxide. Formulae of compouds: Fe2O3 + CO Fe + CO2 Count the number of atoms 2[Fe] + 1[C] 1[Fe] + 1[C] 3 [O] 1[O] 2[O] Balance the iron atoms Fe2O3 + CO 2Fe + CO2 2[Fe] + 1[C] 2[Fe] + 1[C] 3[O] 1[O] 2[O] Balance the oxygen atoms 3[O] 3[O] 6[O] Check the carbon atoms 2[Fe] + 3[C] 2[Fe] + 3[C] 92 Fe2O3 + 3CO 2 Fe + 3CO2 Write the balanced equation when iron reacts with hydrochloric acid to form iron(II) chloride, and hydrogen Write the balanced equation when aluminium hydroxide, decomposes on heating to form aluminium oxide, and water. 93  Problem: What information can be derived from the following chemical equation? 2 CO (g) + O2 (g) → 2 CO2 (g) Solution Mr (CO) = 28 M (CO) = 28 g mol−1 Mr (CO2) = 44 M (CO2) = 44 g mol−1 Mr (O2) = 32 M (O2) = 32 g mol−1 From the equation we can derive that: 2 molecules of CO react with 1 molecule of O2 to yield 2 molecules of CO2 2 moles of CO react with 1 mole of O2 to yield 2 moles of CO2  Cont. 2 CO (g) + O2 (g) → 2 CO2 (g) Substituting the molar masses we obtain: 2 × 28 g of CO react with 32 g of O2 to yield 2 × 44 g of CO2 Assuming that the molar volume of gas: VM = 22.4 L mol−1 (at 0 ºC and 101.3 kPa): 2 × 22.4 L of CO react with 22.4 L of O2 to yield 2 × 22.4 L of CO2 How can we use these statements?  Problem: What volume of oxygen is needed to convert 6 mols of CO to CO2? Derive the result from the previous information It follows the previous expressions 2 mols of CO react with 1 mol O2 that occupy 22.4 L 6 mols of CO will react with 6/2 = 3 mols O2 that occupy 3 x 22.4 L = 67.2 L 6 mols of CO will react with 67.2 L of O2 giving CO2 How many mols of CO2 will be formed? 96 Problem: How many grams of sodium carbonate and calcium chloride are needed to produce 200 g of calcium carbonate? Ar(Na) = 23.0, Ar(C) = 12.0, Ar(O) = 16.0, Ar(Ca) = 40.1, Ar(Cl) = 35.5 Solution Mr (Na2CO3) = 106.0; Mr (CaCl2) = 111.0; Mr (CaCO3) = 100.1 CaCO3 is prepared according to the equation (precipitation reaction): Na2CO3 + CaCl2 → CaCO3 + 2 NaCl 1 mole of Na2CO3 reacts with 1 mole of CaCl2 to yield 1 mol of CaCO3 The amount of CaCO3 needed is n = m/M = 200 (g) / 100.1 (g mol−1) = 1.998 mol The same amount of reactants must be used: m(Na2CO3) = 1.998 (mol) × 106.0 (g mol−1) = 211.8 g 97 m(CaCl2) = 1.998 (mol) × 111.0 (g mol−1) = 221.8 g Problem Calculate the volume of hydrogen (at 0 ºC and 101.3 kPa) which is produced in a reaction of 10.0 g of metallic zinc with an excessive amount of hydrochloric acid. Ar (Zn) = 65.4, Ar (Cl) = 35.5 98 Problem How many grams of silver chloride can be formed when solutions containing 50 g of CaCl2 and 100 g of AgNO3 are mixed? Which of the reactants is in excess and how many grams of this reactant remain unreacted? Ar (N) = 14.0, Ar (O) = 16.0, Ar (Ca) = 40.1, Ar (Cl) = 35.5, Ar (Ag) = 107.9 99 Names and formulas of molecular compounds When naming a molecular compound, the first nonmetal in the formula is named by its element name; the second nonmetal is named using the first sylable of its element name, folowed by -ide. When a subscript indicates two or more atoms of an element, a prefix is shown in front of its name Formula name CO2 Carbon dioxide NO Nitrogen oxide N2O Dinitrogen oxide SF6 Sulfur hexafluoride SiB4 Silicon tetraboride 100 B2O3 Diboron trioxide Complete names and formulae of compounds Formula name Chlorine tetraoxide Phosphorus trichloride Silicon tetrachloride IF3 N2O5 Br2O NO2 SO3 101  Repetition What are Isotops? 102  Isotopes Atoms of the same element that have the same atomic number but differ in number of neutrons The majority of natural elements exist as a mixture of isotopes, one usually predominating over the others Isotopes have the same chemical properties but can differ in some physical properties and have different stability 103 Repetition Oxygen 1) How many protons? How many electrons? 2) Electronic configuration? 3) How many O2 molecules are in 67.2 L? 104 Which atom has the largest atomic size? Mg, Ca, Cl 105 Atomic size 106 Why cations of alkaline metals have always charge 1+? Why cations of alkaline earth metals have always charge 2+? 107  Example: assess the bond polarity in the following compounds: Compound F2 HF LiF Electronegativity difference Type of bond Electronegativities: H – 2,2 Li – 1,0 F – 4,0 108  Suggest structure CO2 NH3 H2 O 109 110 Specify weak interactions HF Br2 111  That is all for today! Thank you for your attention and… Have a nice stay (and further study) in the Czech Republic 112