Grade_11_University_Chemistry_Review

Matter and Bonding 1. Understand the atomic models of Dalton, Thompson and Bohr-Rutherford. a. Dalton Atomic Model: The ‘Billiard Ball’ Model * All matter is made up of tiny particles called atoms. * All atoms of the same element are the same in mass and volume. * Atoms of different elements combine in small whole number ratios to form new compounds. b. Thompson Atomic Model: The ‘Chocolate Chip Cookie’ Model * The discovery of electrons: cathode ray tubes * The atom is a solid ball of positive charge with negatively charged electrons embedded in it. * The amount of positive charge equals the amount of negative charge, so the entire atom is neutral. c. Rutherford Atomic Model: ‘Beehive Model’ * The discovery of the nucleus. * The atom is mostly empty place. * Most of the mass is concentrated in a small, dense nucleus. * Electrons move around the nucleus like planets around the sun. * Discovery using the gold foil experiment: alpha particles are shot into the gold foil. Being that the atom is entirely empty space (prior knowledge), the particles should completely go through the foil. However, some bounced back, repelled. Meaning that the atom is NOT entirely an empty space. The alpha particles are positively charged, therefore, the ‘nucleus’ discovered is also positively charged. * The nucleus’ size is determined by the rarity of the reflected particles. d. Bohr Atomic Model: ‘Energy Level Model’ * Electrons travel around the nucleus in specific energy levels. * The first energy level (orbit) can hold a maximum of 2 electrons. The others can hold a maximum of 8. e. Bohr-Rutherford Model: combination of nucleus and energy levels 2. Outline the name, location, mass and charge of common subatomic particles. Name | Relative Mass | Charge | Location | Electron | 1 | 1- | On energy levels | Proton | 1836.12 | 1+ | In the nucleus | Neutron | 1838.65 | 0 | In the nucleus | 3. Outline trends on the periodic table. a. Atomic Radius: the distance from the centre of the atom to the outermost shell that contains electros ↓ increasing atomic radius | → decreasing atomic radius | Why: number of shells increased | Why: the nuclear charge increases; strength of attraction between nucleus and valence electrons increase as protons are added | b. Ionization Energy: the amount of energy (kJ) that is required to remove one of the outermost electrons from the atom in a gaseous state. The energy to strip an electron from an atom. ↓ decreasing ionization energy | → increasing ionization energy | Why: the bigger the radius, the more energy shells: the further shells have less gravitational pull, and it is easy for this atom to lose its electrons | Why: higher ionization energy atoms have an almost full shell and it is easier to gain electrons; a higher nuclear charge (atoms get smaller): nucleus and valence electrons attractions are increases (more secured valence electrons) | c. Electrongativity: a measure of the ability of the atom to attract an electron from another atom. If an atom is really good at ‘stealing’ electrons from another atom then it has a high electronegativity value. The numbers range from 0.5-4.0 (Fluorine). The ‘Noble Gases’ do not have electronegativity values, since they are already stable. ↓ decreasing elctronegativity | → increasing electronegativity | Why: further shells have less gravitational pull and more likely to lose electrons | 4. Be able to classify compounds as ionic or covalent in nature. Give properties of ionic substances and contrast these with properties of covalent substances. Ionic compounds: a pure substance formed from a metal and a non-metal; metal and polyatomic ion. Held together by attractive forces; positive-negative. * Remember multivalent metals: metals with more than 1 possible charge. Molecular compounds: a pure substance formed from 2 or more different non-metals * Covalent bond: the attractive force between 2 atoms that results when electrons are shared by the atoms; a type of chemical bond * Diatomic molecule: a molecule consisting of 2 atoms of the same or different elements | Molecular (Covalent) | Ionic | Example | Table Sugar | Table Salt | Composition | Made up of molecules | Made up of repeating formula units, which form a crystal lattice. Each ion has a position in the lattice that maximized the attractive forces and minimizes contact with similarly charged ions. | State at SATP | Gas, liquid, or solid | Usually solid | Melting/Boiling Points | LOW | HIGH | Hardness of crystalline solids | Solids are soft and waxy (e.g. candle wax) | Solids are hard and brittle | Solubility in water | Range in solubility (e.g. sugar dissolves; wax doesn’t) | Soluble in water (dissolves) | Conductivity in water | Non-electrolyte | Electrolyte | 5. What is a formula unit' How is the formula unit for an ionic compound determined' Formula unit: the simplest whole-number ratio of atoms or ions of the elements in an ionic compound. The empirical formula of any ionic or covalent network solid compound. A formula unit can be identified by finding the ratio of the ions. NaCl has a 1:1 ratio of sodium and chlorine atoms. The anions and cations in an ionic compound locked in a regular structure, held by the balance of attractive bonds and electrical repulsion. 6. Why are ionic bonds strong' An ionic bond is when one atom gives an electron to another atom. The donor atom gains a positive charge and the receiving atom gains a negative charge. Opposites react and they stick together. (Covalent bonds are stronger) 8. What are covalent bonds' Why do atoms form covalent bonds' What type of atoms tend to form covalent bonds' Covalent bond: the attractive force between 2 atoms that results when electrons are shared by the atoms; a type of chemical bond. To fill their valence shells, two non-metal atoms would share their electrons, so each one has a complete valence shell. 9. How many electrons are shared in a single bond' In a double bond' In a triple bond' Rank single, double and triple bonds in terms of their strength. Triple Bond: three electrons from each atom are shared (T: 6) Double Bond: two electrons from each atom are shared (T: 4) Single Bond: one electron from each atom is shared (T: 2) Bonding Capacity: the number of electrons lost, gained, or shared by an atom when it binds chemically 11. Define bond polarity and use electronegativity values to determine the of bond polarity. Be able to identify the slightly negative and slightly positive ends of bonds. Use the polarity of bonds to predict the polarity of simple molecules. Polar Covalent Bond: a covalent bond formed between atoms with significantly different electronegativities; a bond with some ionic characteristics Polar Molecule: a molecule that is slightly positively charged at one end and slightly negatively charged at the other because of electronegativity differences a. Pure Covalent: 0-0.4 (shared electrons equally) b. Polar Covalent: 0.4-1.7 (polar) c. Ionic: 1.7-3.3 (pulls the electron off) 12. Name simple binary ionic compounds including those with multivalent metal ions. Binary Compound: a compound composed of 2 kinds of atoms or two kinds of monatomic ions Ba2S; MgCl2 Multivalent: the property of having more than 1 possible valence (charge) Tin (III) Oxide 13. Name simple binary covalent compounds. CO2: Carbon Dioxide 14. Name ionic compounds containing polyatomic ions. KNO3: Potassium Nitrate 15. Know four factors that can affect reaction rate. Be able to explain how each factor affects reaction rate. a. Surface Area: more surface area; faster reaction rate * When a solid takes part in a reaction, only the particles at the surface of the solid can react.  If there is more surface, there are more particles that can react and therefore, there are more collisions that can take place, increasing the reaction rate. b. Catalyst(s): does not get consumed, but speeds up reactions c. Temperature: increase temperature; faster reaction * Collision theory: hotter; increase the amount of collisions * ***** d. Concentration: higher concentration; faster reaction rate * A high concentration means that there are more particles in a given volume of space.  This leads to more collisions, which leads to a higher reaction rate. e. Nature of the reactants: what the reactants are changes the reaction rate * There are many factors that affect how a substance will react, such as: a) chemical and physical properties b) state c) molecular structure d) type of bonding e) strength and number of chemical bonds f) intermolecular forces 16. Define, give a general format for, and write out: combustion, synthesis, decomposition, single displacement, and double displacement reactions. a. Synthesis Reaction: a chemical reaction in which 2 or more substances combine to form a more complex substance A+B=AB b. Decomposition Reaction: a chemical reaction in which a compound is broken down into 2 or more simpler substances AB=A+B c. Single Displacement Reaction: the reaction of an element with a compound to produce new element and a new compound A+BC=AC+B d. Double Displacement Reaction: a reaction in which aqueous ionic compounds rearrange cations and anions, resulting in the formation of new compounds AB+CD=AD+CB e. Neutralization Reaction: a double displacement reaction between an acid and a base to produce an ionic compound (a salt) and usually water. MgOH2 s+2 HClaq→MgCl2 aq+2 H2Ol f. Combustion Reaction: the reaction of a substance with oxygen, producing oxides and energy S(s)+O2 (g)→SO2 (g) i. Combustion of Hydrocarbons: Complete Combustion * Excess O2 (g) is present to react * Reaction products are CO2 and H2O(g) * Hydrocarbon combust (burns) with a clean flame ii. Combustion of Hydrocarbons: Incomplete Combustion * Insufficient O2 (g) is present (limiting reagent); there is excess hydrocarbon present. * Reaction products are either CO and/or C and H2O(g) 17. Be able to predict the products of the combustion, synthesis, decomposition, single displacement and double displacement reactions using oxidation numbers. * Single Displacement: The less reactive metal gets displaced by the more reactive metal in a compound (activity series). * Double Displacement: Usually occurs in aqueous solutions (whenever there is a precipitate: use solubility table). Reaction Calculations 3. Define and give examples of a mole. What is Avogadro’s constant' Determine the number of atoms or molecules from the number of moles given. Mole: the amount of substance containing 6.02 x 10^23 entries. For example, 1 mol of people: 6.02 x 10^23 people, 1 mol of oranges: 6.02 x 10^23 orange Avogadro's constant: the number of entities in one mole: 6.02 x 10^23 entries/mol; SI symbol NA. •Avogadro's constant is useful, because there are 6.02 x 10^23 carbon atoms in 12 g of C-12. Each atom of C-12 has a mass of 12 u, and each mole of C-12 atoms (6.02 x 10^23 atoms) has a mass of 12 g. 8. Define the terms limiting and excess reactant. Be able to use these terms to explain the theoretical yield of a reaction. Empirical formulas: the simplest whole number ratio of atoms and ions in a compound. It is usually determined from observation in an experiment and not theory. Empirical formulas are clay layers with % composition ratios. Keep extra decimal places. Molecular formulas: the actual number of atoms in a substance. 9. Be able to balance simple nuclear equations and discuss the four main types of nuclear reactions. Outline the processes of alpha and beta decay and be able to write out equations for each type of decay. a. Alpha (nucleus of a helium atom): variable but relatively slow; a few cm b. Beta (electrons): variable, but relatively fast; a few m c. Gamma: very fast (speed of light); unlimited penetration in the air Alpha Decay: an unstable nucleus that releases alpha particles. This tends to happen with large elements with over 82 protons. An alpha particle has 2 protons and 2 neutrons (Helium nucleus) ZAX→Z-2A-4Y+24α Beta Decay: often a nucleus will have too many neutrons, the nucleus decays by emitting an electron. The extra neutron will break down into a proton, electron, and anti-neutron (v). ZAX→Z+1AY+-10β+v Gamma Decay: often after alpha/beta decay the nucleus is unstable and to attain equilibrium a neutron will release energy. A high energy photon is released. ZAX→ZAX+00γ Beta-negative Decay: an electron is emitted from the nucleus of a parent atom. When a nucleus contains too many neutrons, the strong nuclear force becomes much greater than the electrostatic force. To maintain stability, a neutron spontaneously decays into a proton and an electron, and the electron is ejected from the nucleus Beta-positive Decay: a proton changes into a neutron and a positron, which is symbolized by+10e. Positron is emitted. Gamma Decay: after a nuclear reaction such as alpha or beta decay has occurred, the daughter nucleus is in a high-energy, or excited, state. As a result, the nucleus spontaneously releases energy in the form of a gamma ray in order to return to a lower, more stable energy state. Gamma ray is emitted as a photon. Photon: a high-energy particle with no mass. * The half-life of any given isotope is actually an average time for a particular parent atom to decay to its daughter atom. Solutions and Solubility 1. Define solution, solvent, solute, saturated, unsaturated, supersaturated. Give examples. Solution: a homogenous mixture of substances composed of at least one solute and one solvent. Solvent: the medium in which a solute is dissolved; often the liquid component of a solution. Solute: a substance that is dissolved in a solvent to form a solution. Solubility: a property of a solute; the concentration of a saturated solution of a solute in a solvent at a specific temperature and pressure. Saturated Solution: a solution containing the maximum quantity of a solute at specific temperature and pressure conditions (liquid honey). 2. Define electrolyte and give examples. Electrolyte: a compound that, in aqueous solution, conducts electricity (empirical definition) (e.g Ionic Compunds: NaCl) 3. Outline the process by which an ionic solid is dissolved in water to produce aqueous ions. * IT DISSOCIATES. * Particles in a substance, when dissolving, separate from each other and disperse into the solution. Non electrolytes disperse electrically neutral particles throughout the solution. * A compound such as table salt dissolves, it dissociates into individual aqueous ions. The positive ions are surrounded by negative ends of the polar water molecules, negative ions are surrounded by positive ends of water polar molecules. * Non polar solutes dissolve in non polar solvents. They do not show dipoles or hydrogen bonds. * They have London dispersion forces. Weak intermolecular forces. Explains the relatively low boiling points of non polar and non hydrogen bonding compounds. Similarly, these same compounds dissolve in each other to form solutions. * Water molecule is a universal solvent: small size, highly polar nature, onside able capacity for hydrogen bonding. 4. Outline the concept of “like dissolves like.” Polar substances dissolve in polar substances and non polar solutes in non polar solvents. 5. Outline how the concentration of a solution may be expressed and be able to calculate concentration of solutions. * ppm/ppb -> ppm = 1 mg/L * %v/v * %w/v * %conc. 6. Outline the preparation of a standard solution in a lab. Give all steps. Standard Solution: a solution for which the precise concentration is known. * Start from a solid. * Dissolve in solvent. * Calculate concentration. 7. What is dilution' Be able to do dilution calculations using C1V1= C2V2. Dilution: preparing a solution of lower concentration from a solution of higher concentration. 8. What is solubility' What is a solubility curve. Give definitions of low solubility and high solubility using concentration terminology. Relate patterns of solubility for solids in liquids, gases in liquids and gases in gases. Solubility: a property of a solute; the concentration of a saturated solution of a solute in a solvent at a specific temperature and pressure. Solubility curve: since, the solubility of a substance changes with the temperature of the saturated solution, it is sometimes useful to plot graphs of the relationship between these two variables. The relationship between the solubility and the temperature is graphed. High Solubility: with a maximum concentration at SATP of greater than or equal to 0.1 mol/L. Low Solubility: with a maximum concentration at SATP of less than 0.1 mol/L. * Solids generally have higher solubility at higher temperatures (ionic/molecular) * Gases generally have a higher solubility at colder temperatures (pop in freezer). 9. Write out ionic and net ionic equations for reactions involving solutions. Be able to use a solubility table to predict the formation of precipitates. 10. Be able to use all methods and equations related to solution stoichiometry, especially n=cv. 11. What are acids and bases' Outline the design and organization of the pH scale. What is a logarithmic scale' According to Arrhenius’ Theory Bases: an ionic hydroxide that dissociates in water to produce hydroxide ions. Acids: a compound that ionizes in water to form hydrogen ions. 12. Name indicators and the colours that they appear in acidic and basic conditions. 13. What is a strong versus a weak acid' What is a strong versus a weak base' Give examples of each. 14. Name common binary acids and oxyacids. Name simple bases. 15. Calculate pH from the concentration of hydrogen (hydronium) ions. Be able to reverse this calculation to find [H+]. 16. Outline acid base theories of Arrhenius and Bronsted-Lowry. Name and identify conjugate acid-base pairs. 17. Outline the concept of neutralization. Write out neutralization reactions. 18. Give examples of titration using known concentrations of acid or base.