1413 Chem Notes PDF

Chapter 1 Matter and Measurements 1  Matter Chemistry is the study of matter and the changes that matter undergoes Matter is defined as anything that has a mass and occupies space (volume) and is made up of particles. States of...

Chapter 1 Matter and Measurements 1  Matter Chemistry is the study of matter and the changes that matter undergoes Matter is defined as anything that has a mass and occupies space (volume) and is made up of particles. States of Matter – Solid particles close together in orderly fashion little freedom of motion a solid has a fixed volume and shape – Liquid particles close together but not held rigidly in position particles are free to move past one another a liquid sample has a fixed volume but conforms to the shape of the part of the container it fills – Gas particles randomly spread apart particles have complete freedom of movement a gas sample assumes both shape and volume of container. 2  Classification of Matter Matter is either classified as a pure substance or a mixture of substances. Substance can be either an element or a compound. A mixture can be either homogeneous or heterogeneous. Compounds could be separated to their elements by chemical methods such a heat decomposition and electrolysis. Mixtures could be separated to pure substances by physical methods such as ultrafiltration and distillation Such as ultrafiltration and distillation Such as heat decomposition and electrolysis 3  Pure Substances Element cannot be separated into simpler substances by chemical means. Examples: iron (Fe), mercury (Hg), oxygen (O), magnesium (Mg) and hydrogen (H) Compound is formed when two or more elements chemically combined in definite ratios. Examples: salt (NaCl), water H2O, ethane (C2H6), carbon dioxide CO2, HgO, CaCO3, CaSO4·H2O, Na2CO3, Mg(OH)2 The properties of the compounds are different from the properties of their elements Compounds can be separated into their elements by chemical means such as: Heat decomposition of mercuric oxide (HgO) HgO(sd) Hg(lq) + ½ O2(g) Electrolysis of water : H2O(lq) H2(g) + 1/2O2(g) 4  Mixtures Mixture: physical combination of two or more pure substances – Substances retain distinct identities Types of Mixtures – Homogeneous Mixtures : A homogeneous mixture is one in which the components are uniformly distributed. The composition of the mixture is uniform throughout Example: sugar dissolved in water, air, sea water, NaOH solution (and all solutions), … q Solution is a homogenous mixture of two or more substances that are chemically unreacted. Solute ü Solution is composed of the solute and the solvent. ü Solute: the substance exists in the smallest amount ü Solvent: the substance exists in the largest amount. Solvent Example: in NaCl solution, NaCl (sd) is the solute and water H2O is the solvent. 5  Heterogeneous Mixtures : mixture is made of different substances that remain physically separate. Heterogeneous mixtures always have more than one phase (composition is not uniform throughout) Example: sugar mixed with iron filings, sand and chalk powders in water, water and oil Mixtures can be separated to their components by different physical methods such as filtration and distillation: ØFiltration: used to separate a heterogeneous solid-liquid mixture Example: mixture of water and sand ØDistillation: used to resolve a homogenous solid-liquid mixture Example: salt NaCl in water, sea water 6  Properties of matter q Chemical properties: The ability of a substance to combine with or change into one or more other substances Examples: Heat of combustion, enthalpy of formation, electronegativity q Physical properties: Characteristics that can be observed or measured without changing the composition of the substance Examples: Temperature, color, volume, mass, area, pressure, melting point, boiling point Ø Intensive properties are properties which do not depend on the amount of matter Examples: Temperature, density, boiling point, concentration, solubility, color Ø Extensive properties are properties which depend on the amount of matter Examples: mass, volume, surface (area), amount of substance (number of moles) 7  Chemical Change q A chemical change involves making or breaking chemical bonds to create new substances. Chemical changes include the followings: Oxidation reduction (Redox) reactions: loss and gaining of electrons: The oxidation of iron(II) to iron(III) by hydrogen peroxide in acidic medium Fe2+ Fe3+ + e- H2O2 + 2 e- 2OH- Overall equation: 2Fe2+ + H2O2 + 2H+ 2Fe3+ + 2H2O (Redox reaction) Reaction of base and acid (neutralization reaction): HClsol + NaOHsol H2O lq + NaCl aq Heat decomposition: HgO(sd) Hg(lq) + ½ O2 (g) Electrolysis of water: H2O(Lq) H2(g) + 1/2O2(g) Iron rusting (corrosion): 4Fe + 3O2 2Fe2O3 Combustion (of ethane C2H6): 2C2H6 + 7O2 4CO2 + 6H2O Burning of wood 8  Physical Change q A physical change alters a substance without changing its chemical identity. q No new substance is created, no formation of new chemical bonds during a physical change. q physical changes occur when substances are mixed but don’t chemically react Physical changes include the followings: Phase changes such as vaporization, condensation, freezing, sublimation, melting and deposition. Example: H2O(lq) H2O(g) Dissolving sugar and salts in water Mixing sand with water or mixing oil with water crushing 9  PHYSICAL QUANTITIES and the System International (SI) units q Fundamental (basic) quantities cannot be defined in terms of other physical quantities Examples: quantity Length Mass Temperature Time Amount of substance Examples:, SI basic unit m kg K (Kelvin) s (second) mol (meter) q Derived quantities can be defined in terms of the fundamental physical quantities Examples: quantity area volume density concentration Pressure SI derived m2 m3 kg/m3 mol/m3 Pa (Pascal) unit 10  SI Prefixes Practice: Convert the followings to the desired unit: 23 Gg = …………dg 2 ng = ………mg 2.126 mg = ……..µg Answer: 23 Gg = 23×109 g =23×109×10 dg =23×1010 dg 2 ng = 2×10-9 g = 2×10-9 ×103 mg = 2×10-6 mg 2.126 mg = 2.126×10-3 g = 2.126×10-3×106 µg = 2.126×103 µg = 2126 µg 11 Pressure: P = F/S (F: force and S: surface) The SI unit of pressure is Pascal (Pa) Temperature: – Celsius Represented by °C Based on freezing point of water as 0°C and boiling point of water as 100°C – Kelvin Represented by K (no degree sign) The absolute scale – Fahrenheit (the English system) (°F) The SI unit of temperature is Kelvin (K) Equations for Temperature Conversions K = o C + 273 9 o F = ( ´ o C) + 32 = (1.8 ´ o C) + 32 5 12  Practice on temperature conversions Example 1 : A clock on a local bank reported a temperature reading of 28 oC. What is this temperature in K and °F? Answer: K = 28 o C + 273 = 301 K 9 o o F=( 9 ´ 28 o C) + 32 = 82 o F o F = ( ´ C) + 32 = (1.8 ´ o C) + 32 5 5 Example 2 : A child has a temperature of 104 o F. What is his temperature in (o C) and (K)? Answer: (o C) = [104 -32]/1.8 = 40 o C (K) = 40 + 273 = 313 K Example 3 : A temperature of 300 K, when expressed in °F, would be a) 392 °F. b) 80.6 °F. c) 27 °F. d) 0 °F. 300 K = (300-273) °C =27 °C =(1.8×27)+32 °F= 80.6 °F 13  Volume is an extensive property. Volume is a derived quantity. The SI derived unit of volume is m3 (meter cubed) but the unit liter (L) is more commonly used in the laboratory setting. 1m3 = 1000 L 1dm3 = 1L 1cm3 = 1mL Practice: Convert the followings to the desired unit: a) 1m3 = ………… mL. b) 1 nL= …………. cm3 Answer: a) 1m3 = 1000 L = 103 L = 103 × 103 mL = 106 mL. b) 1 nL= 1 × 10-9 L = 1 × 10-9 ×103 mL = 10-6 mL =10-6 cm3 14  Density: Density is an intensive property (does not depend on the amount of matter). Density is a derived quantity. m d= V d = density; m = mass and V = volume – The SI unit for density is kg/m3 There are other common units for density such as (g/cm3 or g/mL, g/L) (*gas densities are usually expressed in g/L) q Practice 1: A sample of mercury has a volume of 0.25 L and a mass of 3.4 Kg. Calculate the density of mercury in g/cm3? Answer: 1 L = 1000 cm3, Volume = 0.25 x 1000 = 250 cm3 1 kg = 1000 g, Mass = 3.4 x 1000 = 3400 g d = m/v = 3400 g/ 250 cm3 = 13.6 g/cm3 q Practice 2: A cube of metal has a mass of 4.0 g and a length of 2.0 cm on each side. Calculate its density. a) 2.0 g/cm3 b) 0.50 g/cm3 c) 0.20 g/cm3 d) 5.0 g/cm3 15 V =2×2×2 = 8 cm3; d= m/ V = 4/8 = 0.5 g/cm3 q Solubility is the maximum quantity (amount) of the solute that will dissolve in a given amount of the solvent, at a given temperature. Solubility depends on temperature. The solubility may be stated in various units of concentration such as g/L, mol/L and g/100g of solvent. Example: The solubility of sugar in water at 20 o C is 2040 g/L. Solubility is an intensive property Solubility is a derived quantity Practice: The solubility of potassium nitrate (KNO3) in water is 270 g/L at 20 o C. Calculate the mass of KNO3 dissolved in 17.8 mL of water at 20 o C. Ø Answer: at 20 o C, 270g KNO3 1000 mL water (=1L) Mass KNO3 17.8 mL water Mass KNO3 = (17.8 × 270) / 1000 = 4.8 g 16  Uncertainty in Measurement Measurement is the process of comparing an unknown quantity with another quantity of its kind to find out how many times the first includes the second Exact numbers with defined values – Examples: counting numbers, conversion factors based on definitions Inexact numbers obtained by any method other than counting – Examples: measured values in the laboratory SIGNIFICANT FIGURES used to express the uncertainty of inexact numbers obtained by measurement – The last digit in a measured value is an uncertain digit - an estimate 6 or 7 cm 6.7 or 6.8 cm 6.75 or 6.76 cm 17  Rules to determine the significant figures q Any non-zero digit is significant 3455 cm 4 significant figures (4 SF) 28.7 cm 3 significant figures (3 SF) q Zeros between non-zero digits are significant 12051 m 5 significant figures 25400032 nL 8 significant figures 108.036 cm 6 significant figures q Zeros to the left of the first non-zero digit are not significant 0.0058 g 2 significant figures q Zeros to the right of the last non-zero digit are significant if decimal is present 251.00 mL 5 significant figures 30.0 g 3 significant figures 0.002300 dg 4 significant figures q Zeros to the right of the last non-zero digit are not significant if decimal is not present 2500 mL 2 significant figures note that 25.00 mL is a decimal and contains 4 significant figures 18 Practice: Determine the number of significant figures in each measurement. 1234586 cm 7 significant figures 131.21 cm 5 significant figures 0.00018 g 2 significant figures 205 m 3 significant figures 100041 m 6 significant figures 12.00 mL 4 significant figures 3700 mL 2 significant figures 0.0050700 km 5 significant figures 2.145×10-3 m (Scientific notation) 4 significant figures 1.03×104 m (Scientific notation) 3 significant figures 19  Simplified Rounding rules Ø If the number is less than 5 round “down”. 8. 242313 when rounded off to 6 significant figures gives 8. 24231 3. 674 when rounded off to 3 significant figures gives 3. 67 Ø If the number is greater than 5 round “up”. 3.679 when rounded off to 3 significant figures gives 3.68 2.163138 when rounded off to 6 significant figures gives 2.16314 1.5396 when rounded off to 4 significant figures gives 1.540 12.4997when rounded off to 5 significant figures gives 12.500 Ø When the first digit dropped is 5, make the preceding digit even 3.29845 when rounded off to 5 significant figures gives 3.2984 2.13275 when rounded off to 5 significant figures gives 2.1328 3.305 is rounded off to 3.30 if we need 3 significant figures in measurement; Note that zero is an even number 20  Calculations with measured numbers q Addition and subtraction Answer cannot have more digits to the right of the decimal than any of original numbers Example: 102.50 + 0.231 – 12.1 = 90.6 2 digits after 3 digits after 1 digit after 1 digit after decimal point decimal point decimal point decimal point Calculator answer: 90.631 rounded off to 90.6. q Multiplication and division Final answer contains the smallest number of significant figures § Example: 1.4 x 8.011 / 5.12 = 2.2 2SF 4 SF 3 SF 2SF (Limited by 1.4 to 2 significant figures in answer) (Calculator answer: 2.1905078125 round to 2.2) Note that exact numbers do not limit answer because exact numbers have an infinite number of significant figures Example: A penny minted has a mass of 2.5 g. If we have 3 pennies, the total mass is 3 x 2.5 g = 7.5 g In this case, 3 is an exact number and does not limit the number of significant figures in 21 the result. PRACTICE 1: Calculate the following to the correct number of significant figures : (2.20 - 0.0111)/0.1165 = ? Ø (2.20 - 0.0111) = ? (correct answer: 2.19 (2 decimals)) (2 decimals) (4 decimals) (2 decimals) – Calculator answer: 2.1889 2.19 (with 2 decimals) Ø 2.19/0.1165 = ? 3SF 4SF 3SF – Calculator answer: 18.7982832618 18.8 (3 SF) (2.20 - 0.0111)/0.1165 = 18.8 PRACTICE 2: Calculate the following to the correct number of SF: (8.9 × 56) = ………. (2SF) (2SF) (2SF) Calculator answer: 498.4 (4SF) 498.4 = 4.984 ×102 5.0 ×102 (scientific notation) rounded off to 2 SF (8.9 × 56) = 5.0 × 102 (2SF) (2SF) (2SF) 22  Accuracy and precision q Accuracy is how close a measurement is to the accepted value (or true value). q Precision is how closely measurements of the same thing are to one another Practice: The true value of a neutralization volume is 12.370 mL. Describe accuracy and precision for each set of measurements: Set A (12.345 mL, 12.346 mL, 12.344 mL) Precise, not accurate Set B (12.357 mL, 12.337 mL, 12.393 mL) Not precise, not accurate Set C (12.369 mL, 12.370 mL, 12.371 mL) Precise and accurate Answer: The values in set A are close to one another but are not close to the true value so set A is precise but not accurate. The values in set B are not close to one another and are not close to the true value so set B is neither precise nor accurate The values in set C are close to one another and close to the true value so set C is both precise and accurate. 23  Practice on chapter 1 Q1. CaSO4·1/2H2O is a) an element b) a compound c) a heterogeneous mixture d) a homogeneous mixture Q2. The density of water at 20 oC is 0.9982 g/mL. Which of the following sets of density measurements is precise but not accurate? a) (0.9989, 0.9993, 0.9991) b) (0.9981, 0.9980, 0.9982) c) (0.9881, 0.9880, 0.9882) d) (0.8989, 0.8993, 0.8991) Q3. Which of the followings is a derived quantity? a) Mass b) Volume c) Temperature d) Amount of substance Q4. The correct answer to express the arithmetic operation (2.00×0.009)/(9.410-8.0) is a) 0.014 b) 0.01 c) 0.0142 d) 0.0143 Q5. Electrolysis of water is a) a physical change b) a chemical change c) a neutralization process d) none of these Q6. The number of significant figures (SF) in 0.0100 is a) 1 SF. b) 3 SF. c) 4 SF. d) 5 SF. Q7. A temperature of 173 K, when converted to °C equals a) 273 °C b) 446 °C c) -100 °C d) 100 °C Q8. Rounding off 3.1055 to four significant figures results in a) 3.105 b) 3.106 c) 3.1057 d) 3.100 Q9. 600 nm equals a) 6 x 10-3 m b) 0.6 µm c) 6 x 10-2 µm d) 6 x 103 m Q10. Which of the followings is an intensive property? a) Volume b) Mass c) Amount of substance d) Density Q11. Which of the followings is the smallest unit to measure the mass of a substance? a) Mg b) mg c) ng d) µg 24  Chapter 2 Atoms, molecules and Ions 1  The atomic structure Atom consists of three subatomic particles: Protons (p) : have a positive charge (+) Neutrons (n): are neutral and have a mass slightly greater than that of the protons Electrons (e): have a negative charge (-) and a tiny mass Most of the atom is empty Mass of proton ≈ mass of neutron >> mass of electron Atom is the smallest particle (building up unit) of an element, having the same chemical properties as the bulk element. 2  M a ss Num be r A X Elem ent Atom ic Num ber Z Atomic Number (Z) is the number of protons in the nucleus Mass Number (A) is the number of protons and neutrons (P + n) Number of neutrons (n) = Mass Number (A) – Atomic Number (Z) For neutral atom: Number of protons = Number of electrons Examples: M a ss N u m b e r 12 n o. o f (P ) = 6 C n o. o f (e ) = 6 A to m ic N u m b e r 6 n o. o f (n ) = 12 - 6 = 6 197 A u 79 no. of (P) = 79 no. of (e) = 79 no. of (n) = 197 - 79 = 118 3  Isotopes Isotopes are atoms of a given element that differ in the number of neutrons (same atomic number) Examples: Hydrogen 1H 1H 1H Isotopes of hydrogen 1 2 3 H D T Deuterium Tritium Isotopes of carbon Isotopes of oxygen 4  Ions: Anions and Cations A neutral atom becomes an ion either by losing electrons (cation) or by gaining electrons (anion) Na Na+ + 1e- Cl + 1e- Cl- Mg Mg2+ + 2e- O + 2e- O2- Fe Fe3+ + 3e- 5  Molecules A molecule is a neutral group of two or more atoms bonded together that may or may not be the same element. Ø Homonuclear molecules formed by one type of atoms Examples: O2 , O3 , H2 , N2 , Cl2 , Br2 , I2 , F2, … Ø Heteronuclear molecules formed by more than one type of atoms. Examples: H2O, NH3 , NaCl , NaOH, CaSO4·2H2O , Na2CO3 , …. Heteronuclear molecule is the smallest particle (building up unit) of a compound. 6 q Molecular Formula represents the actual number and types of atoms in a molecule Examples: H2 , Cl2 , H2O2 , C6H6 , C6H12O6 , CaCO3, Fe2O3 , CaSO4·2H2O , …. q Empirical or Simplest Formula is the formula that gives only the relative number of atoms of each type in a molecule (gives rise to the smallest set of whole numbers of atoms) Molecular formula = D × (simplest formula) D means how many times simplest formula repeated to give molecular formula. D = 1,2,3… Examples: H2O2 D=2 Empirical formula: HO C6H12O6 D= 6 Empirical formula: CH2O H2O D=1 Empirical formula: H2O C3H6O3 D= 3 Empirical formula: CH2O q Structural Formula is the formula that indicates the attachment of atoms Examples: methane CH4 Ether C2H6O Ethanol C2H6O 7  PRACTICE 1. Which of the followings is the simplest formula of C6H14? a) C6H14 b) CH c) C3H7 d) CHO 2. Which of the followings is the structural formula of methane CH4? a) CH4 b) CH c) C2H8 d) 3. CH2O is the simplest formula of a) C5H1OO b) H-O-H c) CHO d) C3H6O3 4. The molecular formula is the formula that a) gives the relative number of atoms of an element in a molecule b) represents the actual number and types of atoms in a molecule c) indicates the attachment of atoms in a molecule d) represents a group of atoms of the same element containing different number of protons 5. Which of the followings is the building up unit of a compound? a) Atom b) homonuclear molecule c) heteronuclear molecule d) Electron 8  Introduction to the periodic table of elements The arrangement of elements is in order of increasing atomic number (Z) The vertical columns of the table (numbered from 1 to 18 ) are called groups or families. Element in the same group have similar but not identical characteristics. The horizontal rows of the table (numbered from 1 to 7) are called periods. 9  Ionic compounds Most elements are metals. Metals easily lose electrons. Nonmetals tend to gain electrons Metals and nonmetals form compounds with each other. These compounds are called ionic compounds. metal ion + nonmetal ion → ionic compounds q Formula of ionic compounds Ø If the charge of cation = charge of anion, the ionic compound will be 1:1 [cation : anion] Examples: The formula of the ionic compound formed between Na+ and Cl- NaCl Mg2+ and S2- MgS Al3+ and N3- AlN ØIf the charge of the cation and the charge of the anion are different, use the cross over method to determine the formula of the ionic compound. Examples: 10 Q1. 12C, 13C and 14C are called Practice on chapter 2 a) molecules b) compounds c) ions d) isotopes Q2. The formula of the ionic compound formed between Fe and SO42- ions is 2+ a) FeSO4 b) Fe3(SO4)2 c) Fe2(SO4)3 d) Fe3(SO4) Q3. The simplest (empirical) formula of glucose C6H12O6 is a) C6H12 O6 b) C3H4O3 c) CHO d) CH2O Q4. Cations form when a) neutrons are added to or removed from an atom b) electrons are removed from an atom c) electrons are added to an atom d) all of these Q5. The formula that indicates the attachment of atoms in a molecule is the a) structural formula b) ionic formula c) molecular formula d) simplest formula Q6. Which of the following is a homonuclear molecule? a) H2O b) CaF2 c) KNO3 d) Cl2 Q7. The atomic number of an element is the number of a) protons and electrons b) neutrons and electrons c) protons d) protons and neutrons Q8. The numbers of neutrons (n) and electrons (e-) in radium ion 2+ are a) 226 n and 88e- b) 138 n and 88 e- c) 138 n and 86 e- d) 226 n and 86 e- Q9. The horizontal rows in the periodic table are called a) groups or families b) isotopes c) periods d) ionic compounds Q10. The atomic mass of iron (26Fe) is 56. The atomic mass of the ferric ion (Fe3+) is a) 26 b) 56 c) 29 d) 59 3+ Q11. When the cation Ti gains an electron, it becomes a) Ti+ b) Ti4+ c) Ti2+ d) Ti- Q12. The ionic compound is formed between a) two metals b) two nonmetals c) a metal and a nonmetal d) two cations 11