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1 Review Unit 1. Matter a) Elements and Compounds ♣ All matter is made up of about 100 elements. Elements are pure substances that cannot be broken down into simpler parts by ordinary chemical means. Each element is composed of a fundamental particle called an atom. Each element has a unique atom and is represented by a symbol (memorize the sheet “Some Common Symbols”). ♣ The elements with their symbols are organized into the Periodic Table. ♣ Elements combine to form compounds (atoms combine to form molecules). Element 1 (atom) + Element 2 (atom) → Compound (molecule) The physical and chemical properties of a compound are completely different from the elements it is made from. + → two hydrogen atoms + one oxygen atom → one water molecule ♣ Notice that the water molecule can only be made by joining together two hydrogen atoms (symbol = H) with one oxygen atom (symbol = O). The formula for water will be H2O. Note: If there is no number after a symbol in a formula, assume it is a one. Example: CaO means Ca1O1 ♣ In summary, pure substances are organized this way: Matter Elements → Compounds Particles Atoms → Molecules Naming Symbols → Formulas ♣ Impure substances are called mixtures. They are made by mixing different types of compounds and/or elements together. Only mixtures have a variable composition. 1b) Measurement ♣ Measurements are fundamental to the study of matter. There are two types of measurements: (i) quantitative – describes matter using numbers. (ii) qualitative – describes matter in a descriptive, nonnumeric way. ♣ Measurements are reported in significant figures. Significant Figures = all measured digits + one estimated digit (exact) (approximate) 2 43.52 km Example: known uncertain (measured) (estimated) ♣ All of the digits in a measurement are significant unless it is a zero that is used as a place marker. for whole numbers for decimal numbers 6700 0.003 place markers (not significant) place markers (not significant) ♣ To identify significant digits, try this: (i) for whole numbers: 67 490 = count (ii) for decimal numbers: (iii) for mixed numbers: 4 sig. fig. 0.004 80 = count 3 sig. fig. 680.420 = 6 sig. fig. count everything If the zeros in the number were actually measured, they must be shown to be significant. This is done by marking them with a bar or a decimal point. 6700 = 2 sig. fig. (6.7 x 103) 6700 = 3 sig. fig. (6.70 x 103) 6700. = 4 sig. fig. (6.700 x 103) ♣ Remember that significant figures are used only with measurements. They are not applied to defined quantities (100 cm = 1 m) or to pure numbers. ♣ The use of significant figures allows us to change the units of a measurement without changing the accuracy of the measurement. 1.24 m = 1240 mm = 0.00124 km 3 sig. fig. ♣ All measurements use significant figures. When using an instrument: - make sure you understand the scale first. - Remember that the last digit of your measurement must be an estimate (if the spacing is very small, the only practical estimate is either exactly on the line [estimate a 0] or halfway between the lines [estimate a 5].) - be sure the scale is read at eye level. - the level of a liquid is read at the bottom of its meniscus. - Vernier scales can be used to increase the accuracy of the estimated digit. ♣ When significant figures are used in calculations, the answer cannot show more accuracy than the least accurate measurement. It must be rounded off to the proper digit. 3 (1) Addition and Subtraction – round off to the highest place value (the least accurate measurement). 6.4 cm + 8.425 cm + 2.81 cm = 17.635 cm = 17.6 cm the least accurate measurement (tenths) (2) Multiplication and Division – round off to the least number of significant digits used (the least accurate measurement). 6.4 cm x 8.425 cm ÷ 2.81 cm = 19.188 612 1 cm = 19 cm the least accurate measurement ( 2 sig. fig.) 2. Parts of an Atom ♦ The atom has two main parts: (i) the nucleus (or center) of the atom contains positively charged particles called protons (p+) and a neutral particle with no charge called a neutron (n0). It occupies very little space but contains the mass of the atom. (ii) The space in an atom is occupied by negative particles called electrons (e-) traveling around the nucleus at regular intervals. Their have almost no mass. ♦ Atoms are different because their nuclei contain different numbers of protons. Carbon → 6 p+ nitrogen → 7p+ The number of protons in the nucleus of an atom is called its atomic number. Oxygen → atomic number = 8 → has 8 p+ ♦ Almost all of the mass of an atom is in its nucleus. The mass number of an atom, then, is the mass of the particles in the nucleus (protons and neutrons). Oxygen → mass number = 16 → has 16 p+ + n0 A Simplified Lithium Atom = protons = neutrons = electrons atomic number = 3 mass number = 7 atomic symbol = 73Li Nucleus Energy levels ♦ The first 20 electrons are placed in four energy levels ( 2e-, 8e-, 8e-, 2e-) . Their arrangements can be shown using energy level diagrams. Example: show the energy level diagram for potassium (19K) K 19p+ 2e- 8e- 8e- 1e♦ Note: (i) Since atoms have no overall charge, then number of p+ = number of e(ii) number of n0 = mass number – atomic number (iii) Electrons travel around the nucleus in increasing intervals called energy levels. 4 3. Periodic Table ◘ The modern periodic table arranges the elements in order of increasing atomic number. ◘ Metals are separated from nonmetals by the “staircase line”. metals - shiny, malleable, ductile, conductors of heat and electricity. ◘ The columns are families (groups) of elements having similar chemical properties. Some of these families have names to memorize (see sheet). ◘ The rows are called periods. The elements gradually change from metallic to nonmetallic from left to right. The number of the row tells us the number of energy levels for electrons in that atom. ◘ The Group A elements are called representative elements. The Group B elements are called transition elements. ◘ Note that the number given under the symbol is not the mass number of that particular atom. It is an averaged mass number called the atomic mass and will be used later. 4. Ions a) Creating Ions ♣ We know that atoms are electrically neutral because they have equal numbers of protons (p+) and electrons (e-). ♣ When atoms join together, though, they can lose or gain electrons. This causes the charges in the atom to become unbalanced. Ions are atoms that have a charge. atom electron ion metal (M) loses one M+ nonmetal (N) gains one N- ♣ Some atoms will gain or lose more than one electron: - Atoms of metals (M) tend to lose one, two or three electrons and form positive ions (cations). M - 1e- = M+ M - 2e- = M2+ M - 3e- = M3+ Example Mg - 2- = Mg2+ (called a magnesium ion). - Atoms of non-metals (N) tend to gain one, two or three electrons and form negative ions (anions). N + 1e- = NN + 2e- = N2N + 3e- = N3Example O + 2e- = O2- (called an oxide ion). Note that anion names end in “-ide”. 5 4b) Predicting Ionic Charges ▪ It is possible to predict the ion that an atom will form: (i) Representative Elements - these are the Group A elements in the periodic table. The ionic charge for these elements can be determined by their position in the periodic table. 1A M+ 2A M2+ Let M = metal and N = non-metal 3A 4A 3+ transition M metals 5A N3- 6A N2- 7A N- 8A (ii) Transition Elements - most transition (Group B) elements are metals that can form more than one positive ion. iron = Fe2+ and Fe3+ lead = Pb2+ and Pb4+ Since it is difficult to predict these ions, you will be given a list of them to work with (see “Common Ions” sheet). Two methods of naming these ions are used: (a) Stock System - a Roman numeral is used in brackets after the name to indicate the charge. Pb2+ = lead(II) (b) Classical System - the classical name of the element is used with the endings “-ous” (for the lower charge) and “-ic” (for the higher charge). If the symbol is Fe2+ Fe3+ Pb2+ Pb4+ then its name is iron(II) iron(III) lead(II) lead(IV) and its classical name was ferrous ferric plumbous plumbic ▪ Try This Predict the ion that will be formed from these atoms: 1) phosphorus 6) lead(IV) 11) chromium(III) 2) aluminum 7) oxygen 12) calcium 3) cobalt(III) 8) magnesium 13) sulfur 4) nitrogen 9) lithium 14) potassium 5) mercury(II) 10) chlorine 15) tin(IV) 4c) Polyatomic Ions ♠ Poly = many ♠ Polyatomic ions are tightly bound groups of atoms that behave as a unit and carry a charge. Sulfate Ion (SO42-) One sulfur atom and four oxygen atoms forming one strongly bonded unit with two extra electrons. 6 ♠ There are many of these ions. Since we cannot predict their charges, we will be given a list of them to work with (see “Common Ions” sheet). ♠ Most polyatomic groups contain oxygen and their names end in “-ate”. An “-ite” ending means one less oxygen than the “-ate” ending. sulfate = SO42sulfite = SO32- chlorate = ClO3chlorite = ClO2- ♠ A few poyatomic groups do not contain oxygen. Their names end in “-ide”. cyanide = CN5. Formulas a) Creating Formulas ♥ Ionic compounds are created when one kind of positive cation joins with one kind of negative anion. The positive ion is always written first. ♥ The sum of the ionic charges in the formula must be zero. This is done by adjusting the numbers of each ion. Li+ + O2- → Li2O (+1) (-2) (sum=0) Note that the formula for the compound (Li2O) shows no charge. It is neutral overall. ♥How to Balance Ionic Charges - When the charges are opposite numbers, their sum will automatically equal zero. Ca2+ + O2- → CaO (+2) (-2) (sum=0) - When the charges are different, the number of the ions must be adjusted until the sum of the charges becomes zero. Al3+ + Cl- → AlCl3 (+3) (-1) (sum=0) One convenient way of balancing charges is to cross-exchange the digits. Al3+ + O2- → Al3+ 2O2- 3 ♥Note that when you want to show more than one polyatomic group, you must put brackets around the group. Ca2+ + NO3- → Ca(NO3)2 (+2) (-1) (sum=0) 5b) Nomenclature ♦ Let M = Metal N = Nonmetal P = Polyatomic Group H = Hydrogen (1) Ionic Compounds - always begin with a metal - use the “Common Ions” sheet. 7 M + N = M + P = M (stock number) N -ide if necessary M ( ) P -ate (or -ite) NaCl = sodium chloride FeClO3 = iron(III) chlorate (2) Molecular Compounds - no metals present - use prefixes (not the “Common Ions” sheet). N1 + N2 = prefix N1 prefix except mono N2 -ide 1 = mono 2 = di 3 = tri 4 = tetra 5 = penta 6 = hexa 7 = hepta 8 = octa 9 = nona 10 = deca CO = carbon monoxide P2O5 = diphosphorus pentoxide (3) Acids - always begin with hydrogen - there are two types. a) Binary Acids contain only two elements H + N = hydro N -ic acid . HCl = hydrochloric acid b) Oxo-acids contain oxygen in a polyatomic group ite ous H + P ate = P ic acid . H2SO4 = sulfuric acid H2SO3 = sulfurous acid Nomenclature Summary Formula begins with Metal Ionic substance Use the “Common Ions” Sheet Nonmetal Hydrogen Molecular Substance Use Prefixes Acid binary name begins with “hydro-“ oxo“ate-ic ite-ous” acids 8 6. Reactions a) Balancing Equations ♦ A chemical equation is an expression for a chemical reaction. It is a quantitative statement indicating the number of moles of each reactant and of each product. Reactants → Products ♦ In chemical equations, matter must be conserved. The number of atoms of each kind on the reactant side must equal those on the product side. Consider this reaction: carbon + oxygen gas → carbon monoxide C + O2 → CO + However, 2C + O2 → 2CO + → (Matter is not conserved) → (Matter is conserved) This equation is balanced. ♦ All equations must be balanced. Balanced equations have 1. the chemical facts 2. correct formulas 3. atoms conserved ♦ Hints: (1) Try balancing O and H last (2) Never change a formula to make the equation balance. Change coefficients, not subscripts! 2CO2 coefficient subscript (3) Equations that will not balance probably contain an incorrect formula All chemical reaction involves changes to substances. The starting chemicals (reactants) are chemically changed into new chemicals (products) Reactants → Products Chemical equations must obey the Law of Conservation of Matter. Each side of the chemical equation must contain equal amounts of each element. Consider this equation: Carbon + oxygen gas → carbon monoxide gas → CO C + O2 → Since the amounts of each element are not balanced, we use coefficients in front to the formulas to make it balanced. 9 2C + O2 → → 2CO Note: i. the coefficients tell us the amount of each chemical: 2 H2O means 2 mol or 2 molecules of water ii. Equations are balanced only wit coefficients. DO NOT CHANGE THE FORMULAS! Try This: Balance these equations a) H2 + O2 → H2O b) Ca + O2 → CaO c) Mg(ClO3)2 → MgCl2 + O2 d) Ca + H2O → Ca(OH)2 + H2 e) HNO3 + Ba(OH)2 → Ba(NO3)2 + H2O Symbol + → = ⇔ (s) ↓ (l) (aq) (g) ↑ → Pt → heat Symbols used in Equations Explanation Used to separate two reactants or two products “Yields”, separates reactants from products An alternative to → Used in place of a → for reversible reactions Designates a reactant or product in the solid state; placed after the formula Alternative to (s); used only for a solid product (precipitate) Designates a reactant or product in the liquid state; placed after the formula Designates an aqueous solution; the substance is dissolved in water Designates a reactant or product in the gaseous state; placed after the formula Alternative to (g); used only for a gaseous product Indicates that heat is supplied to the reaction A formula written above or below the yield sign indicates its use as a catalyst (in this example, platinum) 6 b) Types of Reactions a) Combination Reactions - Two or more reactants combine to form a single product. R Note: i. Two elements must form a binary compound. S + O2 → SO2 H2 + Cl2 → 2HCl ii. Non-metal oxides in water form an acid. SO3 + H2O → H2SO4 CO2 + H2O → H2CO3 + S → RS 10 Try This: a) C + O2 → b) aluminum + oxygen gas → c) copper (I) + sulfur → d) N2O5 + H2O → b) Decomposition Reactions (Opposite of Combination) - A single reactant is broken down into two or more products: RS → R + S decomposes Note: i. Binary compound 2H2O → 2HgO → into → element + element 2H2 + O2 2Hg + O2 ii. for more complicated compounds, try this: If the compound contains Example C + O NiCO3 H + O NH4NO3 O KClO3 NiCO3 heat → heat NH4NO3 → heat 2KClO3 → One product might be CO2 H2O O2 NiO + CO2 N2O + 2H2O 2KCl + 3O2 Try This: a) H2O → b) mercury (II) oxide heat → c) magnesium carbonate heat → c) Single Replacement Reactions - The atoms of one element can replace the atoms of a different element that is part of an ionic compound. For Metals: M1 + M2X → M1X + M2 Note: that the first metal (M1) can only replace the second metal (M2) if it is more reactive, as listed in the activity series of metals. Examples: i. Ca + 2NaOH → Ca(OH)2 + 2Na ii. Mg + NaOH → no reaction (Mg is not more reactive) For Non-Metals: (mostly the halogen gases) N1 + XN2 → XN1 + N2 Note: that the first non-metal (N1) can only replace the second non-metal (N2) if it is more reactive. Examples: i. Cl2 + BaBr2 → BaCl2 + Br2 ii. I2 + BaBr2 reactive) Try This: a) Fe + CuSO4 → b) Mg + HCl → c) Cl2 + KI → → no reaction (I is not more d) Double Replacement Reactions - This involves the reaction of two compounds to form two new compounds. RS + TU → RU + TS 2KI + Pb(NO3)2 → PbI2 + 2KNO3 Note: i. These reactions involve ionic compounds or acids in aqueous solution. ii. The two reactants just exchange their positive ions. iii. Two compounds will not always react. There are three “driving forces” which cause these reactions to occur: a) a molecular compound can be formed b) a gas can be formed c) a precipitate can be formed - One special type of a double replacement reaction is a neutralization reaction: water Acid + Hydroxide → salt + HCl + NaOH → NaCl + H2O Try This: a) Barium chloride + sodium sulfate → BaCl2 + Na2SO4 → b) Silver nitrate + sodium chromate → 11 Activity Series of the Elements Metals Non-metals Li F Rb Cl K Br Na I Sr S Ba Ca Mg Al MN Zn Cr Fe Cd Co Ni Sn Pb H Sb Bi As Cu Hg Ag Pt Au c) Hydrochloric acid + calcium hydroxide → HCl + Ca(OH)2 → Reactants E1 + E2 C E1 + C1 C1 + C2 SUMMARY Let E= element and C = Compound Reaction Type Products Combination C Decomposition E1 + E2 Single replacement E2 + C2 Double replacement C3 + C4 General Equation R + S → RS RS → R + S T + RS → TS + R RS + TU→RU + TS 12 Other Equation Terms a) Heat - Exothermic reactions are accompanied by the release of heat into their surroundings. CH4 + 2O2 → CO2 + 2H2O + energy The release of energy is shown as a product in the equation. We can show this in an energy level diagram: - Endothermic reactions take place only when heat is continuously to the reactants. CaCO3 + energy → CaO + CO2 The heat is much like a reactant – without it, the reaction cannot take place. Remember that the unit for energy is the Joule (J). b) Physical State - Chemical reactions often depend on the physical state of the chemicals involved. This information can be included in an equation by using these symbols: (s) = solid (l) = liquid (g) = gas (a) = aqueous (dissolved in water) For example, water may be H2O(s), H2O(l) or H2O(g). - Double replacement reactions generally take place between ionic compounds in aqueous solution: PbCl2(aq) + Na2CrO4(aq) These two chemicals will now react if one of the products will form a precipitate that is insoluble in water. PbCl2(aq) + Na2CrO4(aq) → PbCrO4(s) + 2NaCl(aq) The precipitate PbCrO4(s) may now be removed from the solution by filtration. - To write a balanced equation that includes physical states, then, we must be able to predict if a chemical is soluble in water (aq) or insoluble in water (s). To do this, we will use a solubility chart. Note: “s” on a solubility chart = soluble and “(s)” in a equation = solid (insoluble) 13 Try This: Write a balanced equation for each reaction in solution. Include the physical state symbols. 1. Ba(OH)2 + Na3PO4 → 5. (NH4)S + CuCl2 → 2. AgNO3 + K2S → 6. NaOH + CuSO4 → 3. BaCl2 + Na3PO4 → 7. Na2S + Pb(NO3)2 → 4. K2CrO4 + Pb(NO3)2 → 8. PbBr2 + K2CrO4 → Ionic Reactions a) Aqueous Solutions - The particles of a substance, when dissolving, separate from each other and disperse throughout the solution. + water molecule ( ) - When NaCl is dissolved in water, the resulting solution contains ions: NaCl(s) → Na+(aq) + Cl-(aq). The water simply causes the Na+ and Cl- ions to become separated. The process of separating the ions in a solid is called dissociation. - When a crystal contains more than one ion at a particular type, all of the ions will dissociate. CaCl2(s) → Ca2+(aq) + 2Cl-(aq) Al2(SO4)3(s) → 2Al3+(aq) + 3SO42-(aq) - The presence of ions in a solution permits the solution to conduct electricity; the more ions present, the better the ability to conduct. - Many soluble substances, such as HCl(g), are made up of molecules, not ions. These substances, though, create ions when dissolved in water: HCl(g) → H+(aq) + Cl-(aq) This process of splitting a molecule into ions is called ionization. - Precipitation is the reverse of the dissociation process. Ag+(aq) + Cl-(aq) → AgCl(s) b) Ionic Equations A reaction between two chemicals dissolved in water is shown three ways: i. Molecular Equation - the reaction is written as if all of the chemicals consist of molecules: BaCl2(aq) + Na2SO4(aq) → BaSO4(s) + 2 NaCl(aq) ii. Total Ionic Equation - the equation is written to show that the dissolved chemicals are actually in the forms of ions. Ba2+(aq) + 2 Cl-(aq) + 2 Na+(aq) + SO42-(aq) → BaSO4(s) + 2 Na+(aq) + 2 Cl-(aq) 14 iii. Net Ionic Equation - the equation is written to show only the ions or molecules that are actively involved in the reaction. Ba2+(aq) + SO42-(aq) → BaSO4(s) The Na+ and Cl- ions were eliminated since they did not take part in the reaction. Ions like this that occur in the same form on both sides of the equation are called spectator ions. Note: i. Ionic equation should be written from molecular equations. ii. All solution equation must contain the physical state symbols. iii. Ionic Equations are balanced in terms of mass and charge. Try This: Write the total and net ionic equations for the following: a) AgNO3 + K2CO3 → b) Na3PO4 + MgSO4 → c) SnSO4 + Na2S → d) FeBr2 + KOH → → e) SnCl2 + K3PO4