Chemistry Exam Review Grade 11

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Chemistry Exam Review UNIT ONE: PERIODIC TABLE

Group One- alkali metals, most reactive r eactive Group Two- alkaline earth metals Group Three-twelve- are the transition metals, most conduct electricity Group Seventeen- are the halogens, very reactive Group Eighteen- are the noble gases, not very reactive ATOMIC THEORY 

Atomic Number- the number of protons in the nucleus Mass Number- the number of protons and neutrons in the nucleus Number of Neutrons- Mass Number – Atomic Number Isotope- atoms of an element that have the same number of protons but different number of neutrons in their nuclei Relative Atomic Mass (u or amu)- atoms are so small, it is a scale used to indicate average mass. The mass of an atom expressed in the unified atomic mass unit, u. A relative number, compared with the carbon-12 standard. According to th is scale, both the proton & neutron have a mass close to 1u. Example: C-12=12u and H=1 or 1/12 mass of C and O=16 or 16/12 (4/3) mass of C. Average Atomic Mass- is the mass shown in the periodic table that is a result of an average mass of an element based on the abundance of each isotope. Average atomic mass = (atomic mass of A) x (fraction of a) + (atomic mass of B) x (fraction of b)… Radioisotope- a radioactive isotope of an element which is capable of spontaneously emitti emitting ng radia radiatio tion n in the the form form of alpha alpha ( ), beta beta ( ), and and gamm gamma a ( ) rays. rays. They They can occur naturally or can be produced artificially.

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Name Alpha Particle - travels a few cm.

Symbol He or

Beta Particle - travels a few m.

e or

Gamma Rays - penetrating abilities

Description - helium nucleus with no electrons - nucleus lose 2 protons and 2 neutrons, therefore a new element results - high energy e- increases number of  protons by 1 and decreases number of  neutrons by 1 - therefore a new element results - high energy particle of  light or a type of  electromagnetic radiation with no mass/charge - short λ (wavelengths)

Half-life- the time it takes for one half of the nuclei n uclei in a radioactive sample to decay A = A0 (0.5)^t/h A= current sample size, A 0= original sample size, t= time, h= half-life PERIODIC TRENDS

Effective Nuclear Charge- is the net force of attraction between a +ve nucleus and an e- in the valence shell (it tells you how tightly the nucleus is holding onto its e-). - ENC depends on two main factors: magnitude of protons, and electron screening - the magnitude of the nuclear charge (how many p+) does affect the trends across a row (left to right). As you increase the number of p+, ENC increases. - electron screening or electron shielding: the core electrons can affect the force of  attraction between a nucleus and its valence e-. It does affect the trend down a group; the increased number of filled energy levels (core e-) will decrease ENC. Atomic Radius (AR)- magnitude increases across. - nucleus has a stronger force of attraction for it’s valence e-; AR increases from left to right.

Page 3 of 28 - e- shielding causes the nucleus to have a weaker pull on valence e-; AR increases from top to bottom. Ionic Radius (IR)- when discussing ionic radius, we discuss cations (+ve) and anions (-ve) separately. - +ions (cations); removing one e- from an atom causes the electron cloud to shrink.  The extra +ve charge pulls core e- closer. Or a whole outer shell of electrons has been lost (Na+ is smaller than, neutral, Na). - -ve ions (anions); adding an e- to an atom causes the electron could to expand.  The extra –ve charge means e- aren’t pulled closer because there is an increase in repulsion between valence e- (F- is larger than, neutral, F). Ionization Energy- amount of energy required to completely remove an electron from an atom. Electron Affinity- the energy given off when an electron is added. Electronegativity- the electron attracting ability of the atom. NAMING COMPOUDS

If the compound is Ionic (Metal with a Non-Metal)Binary Ionic Compounds: 1. The cation is always written before the anion. 2. The first word of the chemical name is the name of the cation (Ex. Na+ is called Sodium). 3. The last word is the name of the element of which the anion with the suffix – ide. (Ex. Cl- would become chloride). 4. Put the two words together. (Ex. Sodium Chloride). Polyatomic Ionic Compounds: 1. The cation is always written before the anion. 2. The first world of the chemical name is the name of the cation (Ex. Na+ would be called Sodium).

Page 4 of 28 3. The last word is the name of the polyatomic ion (Ex. OH- is hydroxide). 4. Put the two words together. (Ex. Sodium hydroxide).

* For elements in the transition metal group, you must be careful in the naming. Cobalt (II) or (III), Copper (I) or (II), Iron (II) or (III), Lead (II) or (IV), Tin (II) or (IV), Gold (I) or (III) If the compound is Covalent (Non-Metal with Non-Metal)Binary Molecular (covalent) compounds: 1. The first element in the formula uses the whole name of the element, much like binary ionic compounds. (Ex. In NO, the first would be nitrogen). 2. The second element in the formula only uses the first half of the word and – ide is added in place of the removed ending, much like binary ionic compounds. (Ex. In NO, the first part of the second word would be –oxide). 3. So the person reading the name can determine what the subscript is on each element, a prefix is added to show how many of each element are used. The first word of the formula does not use mono. Mono Di Tri Tetra Penta Hexa Hepta Octa Nona Deca * Remember that pairs).

1 2 3 4 5 6 7 8 9 10 certain substance form Diatomic Molecules (they hang out in

H O F Br I N Cl Ion

Name Ammonium Nitrite Nitrate Sulfite Sulfate

Ion

Name Hydroxide Chromate Dichromate Permanganate Acetate

Page 5 of 28 Hydrogen sulfate Peroxide Cyanide Phosphate Hydrogen phosphate Dihydrogen phosphate

Perchlorate Chlorate Chorite Hypochlorite Carbonate Hydrogen carbonate

MOLECULAR OR COVALENT BONDING

Covalent Bonding- a type of bonding which results when one or more pairs of electrons are shared between two atoms (non-metals). - a molecule is the smallest part of an element or compound which exists independently. It contains atoms bonded together in a fixed whole number ratio. - uses prefixes: mono, di, tri… - Cl + Cl : these two atoms have equal IE, EN, Ar, EA. Properties of Covalent Compounds1. Generally have much lower melting and boiling points than ionic compounds. Covalent compounds do not need to break any bonds like ionic compounds. Have weak intermolecular forces which are easy to break/be separated. 2. Are soft, waxy, or flexible. Covalent compounds are free to move around unlike ionic compounds. Do not contain the crystal lattice structure of ionic compounds. 3. Aren’t very soluble in water. Like dissolves like, compounds tend to dissolve in other compounds that have similar properties. Since H 2O is polar and most covalent compounds are not, they don’t usually dissolve in H 2O. 4. They do not conduct electricity in water. Since they do not usually dissolve in water, to form ions they cannot be charge carriers required to conduct electricity. No ions, no electricity (non-electrolyte). BONDING

Bonding- a chemical bond is a strong interaction between atoms. Atoms form bonds to decrease their potential energy which makes the system more stable. Bond type depends on electronegativity.

Page 6 of 28 Ionic Bonds- is a type of chemical bonding resulting from the electrostatic attraction between oppositely charged ions in a compound. The family is called “salts.” Properties of Ionic Compounds- form crystals. Are solids at SATP. - have high melting and boiling points. Most ionic compounds cannot be melted with a Bunsen burner.

- very hard and brittle. Ions resist movement. - are electrolytes. When an ionic compound is put into water the ions pull apart from each other and the negatively charged ions are free to conduct electricity. Lewis Dot Diagrams- determine the number of valence electrons. - write down the symbol (represents nucleus) - fill in the dots (e-) according to the number of valence electrons. Must be in a clockwise fashion. Lewis StructureFor ionic compounds, 1. Determine the Lewis dot diagram for both elements. 2. Transfer the electrons so that both atoms are isoelectric with noble gas. 3. Place brackets around each (always remember to write the cation first, anion anion second) and add subscripts if necessary. Ex. Sodium Chloride

LEWIS THEORY OF BONDING

- atoms and ions are stable if they have noble gas-like electron structure (octet). - e- are most stable when they are paired. - atoms form chemical bonds to achieve a stable octet of e-.

Page 7 of 28 - can achieve stability through ionic or covalent bonding.

Rules for Drawing Lewis Structure1. Arrange atoms symmetrically around the central atom. 2. Count the number of valence electrons of all atoms, add or subtract charges if necessary. 3. Place a bonding pair of electrons between the central atom and each of the surrounding atoms. 4. Complete the octets of the surrounding atoms using one pair of electrons. Any remaining atoms go to the central atom (usually if central atom is period 3 or later, octet rule may not apply). 5. If the central atom does not have an octet, move lone pairs from the surrounding atoms to form double or triple bonds until the central atom has a complete octet. 6. Calculate formal charge and reduce, making resonance structures if  necessary. Formal Charge = # valence electrons – (1/2 #bonded e- + # lone pairs) 7. Draw the Lewis Structure and enclose polyatomic ions with square brackets showing the ion charge. VSEPR SHAPES

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VSEPR THEORY 

- VSEPR Theory provides a 3-D geometric structure to be determined using both the bonded and unbounded electrons of the central atom of a molecule.

Some Key Steps- draw the Lewis Structure for the molecule - once the Lewis Structure is drawn, determine the total number of the bonded and lone pairs of the central atom (steric number) - this total will provide you with one of the basic geometric shapes/molecular shapes - arrange and compare the number of lone pairs with the number of bonded pairs to get the final geometric shape

POLAR COVALENT BONDS

- In many molecules, bonding is neither fully ionic or fully covalent. Ex. H + Cl



HCl

Page 9 of 28 - The molecule as a whole is electrically neutral, however, Cl has a stronger pull of  e- over H. Thus, the shared e- will be closer on average to the Cl atom. - Electronegativity: a measure of the electron attracting ability of the atoms in a molecule. - EN values are such that the bigger its number, the greater that atom’s ability to attract electrons will be. - EN can be used to identify the character of a bond: ionic, covalent, polar covalent. ElectronegativitiesBy determining the difference in EN between 2 atoms, we can determine what type of bond is formed. Covalent

Polar Covalent

Ionic

--------------------------------------------------------------------------------0 Cl + H

1.7 

4.0

ClH

H—EN of 2.1 Cl—EN of 2.9 Difference is 0.8, therefore polar covalent bond. - Cl has a stronger pull on the shared e-, it does not attract strong enough to gain complete possession of it. - therefore the bond is not ionic. - this is known as a polar covalent bond. - the unequal sharing of the bonded e- pair results in a compound where one end is slighty positive ( and the other end is slightly negative ( . - it is almost as if Cl has gained an electron and become negatively charged while H has lost an e- to become positively charged. - when determining whether a molecule is Polar or Non-Polar, we must consider the following: - polarity of each bond in the molecule - bond/atom arrangement in a molecule (VSEPR shape) INTERMOLECULAR FORCES

- are forces of attraction between covalently bonded molecules - three types of intermolecular forces include: London dispersion forces, dipoledipole forces, and hydrogen bonding. - London dispersion forces result from instantaneous dipoles

Page 10 of 28 - occurs from an attractive force acting between all molecules, including nonpolar ones. - very weak forces - they increase as the number of atoms in molecule increases, and as the size of the atom increases - higher boiling point = stronger intermolecular forces - Dipole-Dipole forces result when polar molecules attract nearby polar molecules. - occurs from an attractive force acting between polar molecules. - these molecules also have London dispersion forces. - stronger than London Dispersion forces - Hydrogen bonding - results when an H atom is bonded to a small, electronegative atom - bonded to N, O, or F POLYATOMIC IONS - 4 combinations of the polyatomic ions can exist. In this case we’ll use ClO 3- ion as an example.

- ___________ate oxyanion is the most stable. - Per______ate oxyanion has 1 more oxygen than the ate. - __________ite oxyanion has 1 less than the ate. - Hype________ite oxyanion has 1 fewer than the ite.  Therefore, - ClO- is the hypechlorite ion. - ClO2- is the chlorite ion. - ClO3- is the chlorate ion



most stable.

- ClO4- is the perchlorate ion. * the charge does not change, only the number of oxygen atoms. If you have a polyatomic ion and you add a hydrogen atom, the charge on the ion does not change. If you remove one oxygen atom in exchange for a sulfur atom it becomes  Thio_______ 

HYDRATES

- are compounds that contain water as part of the ionic crystal structure - these compounds decompose to an ionic compound and water vapour when heated

Page 11 of 28 - when aqueous solutions of many soluble salts area evaporated a precise number of water molecules may be retained as the ions for crystals - water becomes a part of the crystal called water of hydration knows as a hydrate How to Name Hydrates1. Name the ionic compound 2. Followed by a dot 3. Prefix the number of water molecules and end in hydrate

Ex. Copper (II) Sulfate Pentahydrate



CuSO4

.

5H2O

NOMENCLATURE CONTINUED

Classic MethodRULE: hydro + stem of element + ic then acid HCL



hydrochloric acid

IUPAC MethodRULE: aqueous hydrogen stem + ide HCL



aqueous hydrogen chloride

Oxyacids: ions that form when hydrogen combines with polyatomic ions containing oxygen

Classic MethodRULE: anion + ous or ic then acid If the ion ends in ate, the acid ends in ic If the ion ends in ite, the acid ends in ous Stem

Ion

Acid Name (add H+ in formula)

1 more oxygen

Per_____ate

Per______ic acid

Stem (most stable)

_______ate

________ic acid

1 less oxygen

________ite

________ous acid

2 less oxygen

Hypo_______ite

Hypo______ous acid

Classic Method ExamplesHCLO4 HClO





perchlorate

hypchlorite





perchloric acid

hypochlorous acid

IUPAC Method ExamplesH2SO4  aqueous hydrogen sulfate HNO4  aqueous hydrogen pernitrate

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CHEMICAL BONDING

Kinetic Molecular Theory (KMT)- explains that chemical reactions occur because particles collide with sufficient energy to break the old bonds and allow the formation of new bonds Collision-Reaction Theory- is a theory stating that chemical reactions involve collisions and rearrangements of atoms or groups of atoms, and that the outcome of collisions depends on the energy and orientation of collisions Representing a Chemical Change- these include a word equation and a chemical equation Evidence of a Chemical Reaction- change in temperature - evolution of a new gas (effervescence) - change in colour - formation of a precipitate (solid produced from two liquids) - formation of an odour - production of light - production of sound  Types of Chemical Reactions- combustion, synthesis, decomposition, single displacement, double displacement Combustion- a very rapid reaction of a substance with oxygen to: - produce compounds called oxides - this type of reaction is also known as burning Complete Combustion- results in the production of water and carbon dioxide - this will only occur if sufficient oxygen is present - General formula: hydrocarbon + oxygen



carbon dioxide + water

Incomplete Combustion- occurs when there is not enough oxygen available - instead of 2 products, 4 are produced: carbon monoxide, carbon, carbon dioxide, and water Synthesis Reactions-A+B



AB

Decomposition Reactions-

Page 13 of 28 - AB



A+B

Single Displacement Reactions- A + BC



AC + B

- in the activity series, each metal will displace any metal listed below it Double Displacement Reactions- AB + CD = AD +CB - use the solubility table UNIT TWO: CALCULATING # OF ENTITES

- used to calculate the # of atoms or molecules in a sample N = nNA - where N = # of entities (atoms or molecules) - where n = moles of sample - where NA = Avogadro’s Constant (6.02 x 10^23 entities) LAW OF DEFINITE PROPORTIONS

- from Dalton’s atomic theory - a specific compound always contains the same elements in definite proportions by mass - For example: CO and CO 2 - the chemical formula for CO is similar to that of CO 2; however, the difference in the number of oxygen atoms per molecule causes each gas to have different properties - even though CO and CO 2 are similar, they have very different properties - properties of a substance depend on the elements the compound is composed of  - we refer to this study of relationships between quantities of reactants and products as stoichiometry THE MOLE

- 6.02 x 10^23 - chemists often need to measure out a certain mass of a substance - use the unit of measurement called the mole (mol) - a mole is Avogadro’s constant - the mass of one mole is called Molar Mass

Page 14 of 28 CALCULATIONS INVOLVING MOLE

- to convert mass to moles, we use the following formula: n=

PERCENTAGE COMPOSITION

- is a calculation used to determine the contribution of each element (as a %) to the total mass of the compound - to determine the percentage composition of a compound, you must: - calculate the mass (or molar mass) of each atom/element in the compound - determine the total mass (or molar mass) of the compound - divide the mass (or molar mass) of each atom/element by the total mass of  the compound to determine the percentage composition of each element/atom in the compound Ex. What is the percentage composition of nitric acid (HNO 3) ? M = 63.02 g/mol %H =

x 100% = 1.6%

%N =

x 100% = 22.23%

%O =

x 100% = 76.16%

EMPIRICAL FORMULA

- indicates a numerical ratio that exists between atoms - the simplest formula/empirical formula does not always indicate the action number of each type of atom within a compound - How to find empirical formula? - Percent to mass - Mass to mole - Divide by small - Multiply till whole

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MOLECULAR FORMULA

- indicates the number and type of each atom found in the molecule - we cannot write a molecular formula for all compounds because not all substances are molecules - in some cases, the empirical formula and molecular formula can be the same - What is the molecular formula?

THE MOLE AND CHEMICAL EQUATIONS

Fe(s) + Se(s)  FeSe (s) - this says: 1 atom of Fe reacts with 1 atom of Se to produce 1 molecule of FeSe - or: 1 mole of Fe reacts with 1 mole of Se to produce 1 mole of FeSe - or: 55.85g of Fe reacts with 78.9g of Se to produce 134.81g of FeSe - But the order from Boreal has not arrived yet. How much of Se is required to react with 16.2g of Fe?  To Solve: - Write a balanced chemical equation (with states) - Record the given mass below that - Record the molar masses below that - Calculate the number of moles LIMITING AND EXCESS REAGENTS

- the reactant that is completely consumed in a chemical reaction is known as the limiting reagent - it determines the around of product that will come from a reaction - if the limiting reagent is all used up, the reaction will stop, leaving an excess of the other reactants  To Solve: - Write a balanced chemical equation - Calculate the number of moles present - Determine the limiting reagent - Determine mass of product PERCENTAGE YIELD

- so far, we’ve been calculating theoretical yield obtained from stoichiometric calculations using our balanced chemical equations

Page 16 of 28 - however, in practice, we don’t always get that expected amount. The amount that we measure at the end of our experiment is the actual yield. - percentage yield compares the actual and theoretical yield. Percentage Yield =

x 100%

BALANCING NUCLEAR EQUATIONS

Nuclear reactions involve changes occurring in the nucleus of atoms - involve the greatest quantities of energy - involve nucleons (particles in nucleus)  There are four classes of Nuclear Reactions: 1) Radioactive Decay - the spontaneous decomposition of a nucleus. Can be wither alpha or beta decay depending on the particle emitted. - an alpha particle is the nucleus of a He atom (no electron) - changing the number of protons changes the atom’s identity. This change is known as transmutation. Ex.

+



Beta Decay: - unstable nuclei may emit high energy electrons known as beta particles in the process of beta decay. +



-

- the electron emitted was never found in the nucleus, but was created by the nuclear reaction: 

+

-

- positrons (positive electrons) are produced by the reverse reaction - both alpha and beta decay produce s more stable nucleus - as nucleons reorganize themselves into a more stable configuration, energy is emitted in the form of gamma rays 2) Artificial Transmutation - is done by bombarding a nucleus with an alpha particle (the nucleus), a proton (hydrogen nucleus) or a neutron Ex. Write a nuclear equation for the bombardment of nitrogen-14 by an alpha particle giving off a proton and the desired isotope. +



+

3) Fission - occurs when an isotope absorbs a neutron, causing it to split into two smaller nuclei plus 2 or 3 neutrons - releases a huge amount of energy

Page 17 of 28 - production of neutrons leads to further fission, the cycle continues - provide energy for nuclear power generating stations 4) Fusion - occurs when several light atoms combine to form heavier atoms to release a very large amount of energy UNIT THREE: THE NATURE AND PROPERTIES OF SOLUTIONS

Solutions are homogeneous mixtures of a substance composed of at least one solute and one solvent. - liquid and gas solutions are transparent - uniform mixture of particles - solutes and solvents may be solids, liquids, or gases Aqueous Solutions: - have water as the solvent (“universal”) - water makes a great solvent because: - a small molecule - highly polar - capable of hydrogen bonding - a liquid over a wide temperature range Intermolecular Forces: - forces BETWEEN molecules define their solubility LONDON DISPERSION (WEAK) - electron in one molecule is attracted to a positive nuclei of others - all (polar and nonpolar) DIPOLE-DIPOLE - between oppositely charged ends - polar molecules HYDROGEN BOND (STRONG) - between slightly positive H atom of one molecule and highly electronegative atom (F, O, or N) of another - polar molecules containing F, O, or N “Like dissolves like”: - polar solutes dissolve in polar solvents due to dipole-dipole and/or hydrogen bonding - non-polar solutes dissolve in non-polar solvents due to London-dispersion forces

Page 18 of 28 Electrolytes vs. Non-Electrolytes: - compounds are electrolytes if their aqueous solutions conduct electricity - most highly soluble compounds - non-electrolytes are compounds that do not conduct electricity in aqueous solution - includes most molecular compounds with the exception of acids Classifying Solutions: Acids, Bases, and Neutral Substances: - acids in aqueous solutions will turn blue litmus paper red - bases in aqueous solution will turn red litmus paper blue - neutral substances will show no change in litmus paper SOLUTION CONCENTRATION

- the ratio between the amount of solute and the amount of solution - in general, Concentration = - dilute solutions have a relatively small amount of solvent per unit of volume of  solution - concentrated solutions have a relatively large amount of solute per unit volume of  solution - concentrations can be expressed in various ways depending on the usage of the solution 1) Percentage Concentration: can be found on many consumer products and may be expressed as a percentage by volume, a percentage by weight or a percentage.

Percentage volume by volume: c=

x 100%

Percentage weight by volume: c=

x 100%

Percentage weight by weight: c=

x 100%

Page 19 of 28 - very low concentrations are often better described in terms of parts per million (ppm) 1ppm = 1g/10^6mL = 1g/1000L = 1mg/L = 1ug/g Molar Concentration: - the amount of solute in moles that is dissolved in one litre of solution ,with units of  mol/L c=

OR M =

SOLUTION PREPARATION

- a standard solution has a precisely known concentration because it was made with a precise mass of solute and a precise volume of solvent - aqueous standard solutions are made using distilled (pure) water as the solvent Preparing a Standard Solution: - standard/stock solutions often have a high concentration, and may not be appropriate for the required purpose - to obtain the desired (lower) concentration we can dilute the solution by adding more solvent - in doing so, the amount of dissolved solute remains the same, while the volume and concentration are changed from the initial solution to the final solution For our final and initial solutions, = SOLUBILITY AND REACTIONS

Saturated Solutions- have maximum solute concentrations- no more solute will dissolve. The solution will remain the same, as long as temperate and pressure are constant. Solubility- is the maximum concentration of a solute in a solvent at a specific temperate (in grams of solute/100mL of solvent). Solubility Curves: - compounds have different solubility’s at different temperatures - solids generally have a higher solubility at higher temperatures - under the curve = unsaturated - along the curve = saturated - above the curve = supersaturated

Page 20 of 28 - gases dissolve in liquids too - gases always have higher solubility at lower temperatures - liquids that do not dissolve in water (mostly non-polar) will form a separate layer and are called immiscible with water - some liquids (mostly polar) dissolve completely in water in any proportions and are said to be miscible with water Crystillization: - occurs when solution concentration exceeds solubility - loss of solvent (through evaporation) increases solute concentration, crystals come out of solution - when concentration increases, volume decreases and moles increase IONIC EQUATIONS

- total ionic equations show all entities present in the reaction with all ionic compounds in their dissociated form. Precipitates have a low solubility in water and are shown as solids. - net ionic equations show only the ions or neutral substances involved in the overall chemical reaction. Qualitative Chemical Analysis: - describes a quality or change in matter that has no numerical expression - some examples: colour, smell, formation of solid, effervescence, light produced, etc. - qualitative analysis by colour: most ions in aqueous solution are colourless, but some have specific colours in solution - chemists use sequential qualitative chemical analysis to test whether ions are present in a sample solution - use formation of low-solubility compounds (precipitate) to detect the presence of ions

Ex. silver in

How can we determine the presence of  and/or lead ions solution?

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SOLUTION STOICHIOMETRY 

- is a method of calculating the concentration of substances in a chemical reaction by measuring the volumes of solutions that react completely; sometimes called volumetric stoichiometry THE DISSOCIATION OF WATER

- water is considered a non-electrolyte - water consists of molecules, not ions - water is actually a weak electrolyte, there are a small number of ions present in water because of dissociation H2O(l) + H2O(l)  H3O(aq) (hydronium ion) + OH (hydroxide ion) - when two water molecules collide, a hydrogen ion is transferred from one water molecule to the other (fewer than two in one billion ionize at SATP)  The Arrhenius Definition: - according to Arrhenius, ionic compounds separate into ions when they are liquid or in solution ACIDS - in their pure form (liquid/solid) acids are molecular compounds that contain hydrogen - they do not conduct electricity and are termed non-electrolytes - however, in solution, acids ionize to release hydrogen ions - hydrogen ions do not exist in water, H 3O or hydronium ions do exist - acids turn litmus paper blue



red

BASES - most bases are ionic compounds that contain a hydroxide, which is released in solution - we can define a base as an ionic hydroxide that releases mobile hydroxide ions in solution - bases turn litmus paper red



blue

Page 22 of 28

STRONG ACIDS - an acid that ionizes almost completely in water (>99%) to form aqueous hydrogen ions (hydronium ions) - a high percentage of ionization, high conductivity WEAK ACIDS - an acid that ionizes only partially (<50%) in water to form aqueous hydrogen ions - low percentage ionization, low conductivity STRONG BASE - a base that dissociates totally - metal hydroxides are strong bases WEAK BASE - a base that dissociates partially THE pH SCALE

- the pH scale is used to measure the strength of acids and bases. 0.0 ---------------------ACID-----------------------7.0---------------------BASE-----------------14.0 - pH is related to the quantity of H+ ions in the solution pH = -log [H+] and [H+] = 10^-pH - ** the number of digits following the decimal point in the pH value is equal to the number of sig. digs in the hydrogen ion concentration** Ex. An antacid solution has a hydrogen ion concentration of 4.7 x 10^-11 mol/L. What is the pH? = -log[H+] = -log[4.7 x 10^-11 mol/L] = 10.328 pH is 10.33 Ex. The pH reading of a solution is 10.33. What is the hydronium ion concentration? [H+] = 10^-pH = 10^-10.33 = 4.7 x 10^-11 mol/L - pH can be determined using a pH meter or an indicator Litmus: red 6.0 – 8.0 blue Phenolphthalein: colourless 8.0 – 10.0 pink Demo: universal indicator, ROYGBV - What effect does the dilution of an acidic solution have on the pH of the solution? Demo: 1.0mol/L HCl, pH = 1

Page 23 of 28 1.0mL of 1.0 mol/L HCl diluted to 100mL, pH = 2 Neutralization Reactions: - when an acid reacts with a base, the pH moves closer to 7 (neutral) Acid + Base



Salt + Water

BRONSTED-LOWRY DEFINITION

- Bronsted-Lowry Acid: a substance that can donate a proton to some other substance. An acid is a proton donor (H+). - Bronsted-Lowry Base: a substance that can accept a proton from another substance. A base is a proton acceptor. - a substance formed by the addition of a proton to a base is called a conjugate acid - a substance formed by the loss of a proton is called a conjugate base - a conjugate acid-base pair may be defined as an acid and corresponding base that differ in chemical composition by only a single proton (H+). Example: NH3/NH4 NH3 (l) + H2O(l)  NH4(aq)+ OH(aq) Base

Acid

C.A.

C.B

 Therefore, H2O/OH and NH3/NH4 are the acid-base pairs. Acid-Base Reactions: - acids can take part in several characteristic reactions that can allow us to get clues about unknown substances or predicting what the products could be 1. Active metal + acid



2. Acid + carbonate

H2CO3(aq) + ionic compound



H2(g) + ionic compound

3. Acid + ionic compound 4. Acid + base





precipitate + acid

salt + water

ACID-BASE TITRATION

- a titration is the procedure where a standardized substance of known concentration (the titrant in buret) is added to a substance of unknown concentration (the sample in Erlenmeyer flask) in an attempt to determine the unknown concentration - usually occurs with neutralization reactions where one of the standardized samples is the base, and the sample is the acid Ex. [HCL] = unknown, [NaOH] = 1.000 M - to determine the [HCl], take a specific amount of the solution (20.0mL) in an Erlenmeyer flask - add 2 drops of indicator solution (phenolphthalein)

Page 24 of 28 - slowly add [NaOH] while carefully measuring the amount with the buret until all of  the HCl has reacted - [H+] = [OH-], this is the equivalence point - the point at which stoichiometry equivalent quantities of the substances have been brought together (equal number of moles) - in the case of acid/base titrations, indicators can be used - Endpoint: occurs when the indicator changes colour and the titration stopped. - ideally, endpoint = equivalence point UNIT FOUR: KINETIC MOLECULAR THEORY (KMT)

- gases consist of extremely small particles called molecules. These molecules are so small that their volume is negligible in comparison with the volume of the container. - the molecules of a gas are in rapid, random, straight-line motion. They collide with each other and with the walls of the container. - all collisions are perfectly elastic; that is, there are no energy losses due to friction. - there are no attractive forces between the molecules. - molecules of different gases have equal average kinetic energies at the same temperature. If the temperature increases, the average kinetic energy of the molecules increases. (Recall: T measures average kinetic energy). Pressure of Gases: - pressure is defined as a force per unit area - atmospheric pressure is the force per unit area exerted by air on all objects - the SI unit for pressure is Pascal (Pa) - atmospheric pressure is approx. 100kPa UNIT NAME

UNIT SYMBOL

DEFINITION/CONVERSION

Pascal

Pa

1 Pa = 1N/m^2

Atmosphere

atm

1 atm = 101.325kPa

Millimeters of mercury

mmHg

760mmHg = 1 atm

Torr

torr

1 torr = 1mmHg

** 101.325 kPa = 1.00atm, 760 torr, 760 mmHg **  Temperature for Gases: - temperature measures the average kinetic energy of a substance’s particles - the common unit for temperature is degrees Celsius, but the SI unit is Kelvin (K) - always use Kelvin when working with gases

Page 25 of 28 Kelvin Scale: - Lord Kelvin found that regardless of the gas tested, the x- on a graph would always be -273 degrees Celsius - molecular motion would cease and so this point is called absolute zero -273 degrees Celsius = 0.0 K = absolute zero K = degrees Celsius + 273 Degrees Celsius = K – 273 - standard temperature and pressure (STP) is 273.15K and 101.325kPa - standard ambient temperature and pressure (SATP) is 298.15K and 100kPa GAS LAWS

Charles’ Law: - the average kinetic energy of gas molecules is directly related to the temperature.  The greater the temperature, the greater the average kinetic energy. - the volume of a fixed mass of gas is proportional to its temperature when the pressure is kept constant.

Gay Lussac’s Law: - Joseph Gay-Lussac discovered the relationship between temperature and pressure acting on a fixed volume of gas. - the pressure of a fixed amount of gas, at constant volume, is directly proportional to its Kelvin temperature.

Boyle’s Law: - Robert Boyle studied the effect of changing the pressure of a gas on its volume at constant temperature. - he determined that as the pressure on a gas increases, the volume of the gas decreases proportionally provided the temperature and amount of gas stayed constant. - at constant temperature, the volume of a given sample of gas is inversely proportional to its pressure. P1V1 = P2V2 Combined Gas Law: - one can see that volume and pressure are directly related to temperature, and inversely related to each other

Page 26 of 28 IDEAL GAS LAW

Purpose: to calculate the amount of gas at any specific conditions of pressure, volume, and temperature Ideal gas: a hypothetical gas that obeys all the gas laws perfectly under all conditions (ex. Does not condense into a liquid when cooled) - the ideal gas law works for any sample of gas - assume that gases behave “ideally” (obeys gas laws under all temperatures and pressures) - ideal gas does not really exist, but makes a close approximation - particles have no volume - no attractive forces - there are no gases for which this is true, although real gases behave this way at high temperatures and low pressures - an equation that gives the relationship between the pressure, volume, temperature and number of moles of a gas is: PV = nRT - where P = kPa - where V = Litres - where n = moles - where R = - where T = Kelvin LAW OF COMBINING VOLUMES

- also known as Gay-Lusaac’s Law of Combing Volumes- when measure at the same temperature and pressure, volumes of gaseous reactants and products of chemical reactions will always combine in simple whole number ratios - Avogadro’s Theory- states that equal volumes of gases measured at the same temperature and pressure contain equal numbers of molecules (and therefore moles) - this theoretical concept helps explain the laws of combining volumes - in a balanced chemical equation, the mole ratio can be expressed as the volume ratio Molar Volume of Gases: - Avogadro’s theory showed that equal volumes of gases contain an equal number of molecules under identical conditions - one mole (6.02 x 10^23 molecules) of a gas should occupy some definite volume under specific conditions - we know two specified conditions of temperature and pressure. They are STP and SATP

Page 27 of 28 - using the ideal gas law, we can determine the volume of one mole of a gas at STP (273.15K and 101.325kPa). This will determine the molar volume of a gas. - One mole of gas occupies 22.4L at STP and 24.8L at SATP - we can use this knowledge to calculate the volume of gases involved in chemical reactions or to calculate the numbers of moles of gases involved in chemical reactions - if volumes aren’t STP or SATP, they will be corrected to STP or SATP DALTON’S THEORY OF PARTIAL PRESSURE

- John Dalton hypothesized that gas particles behaved independently and that the pressure exerted by an individual has is the same whether it exists by itself or as a mixture - conducted a variety of experiments to conclude that each component of air contributes to the total air pressure Ptotal = P1 + P2 + P3… Using KMT to explain Dalton’s Law of Partial Pressures: - pressure of a gas is caused by collisions of molecules with the walls of a container - gas molecules act independently of each other - therefore, total pressure (total collisions with walls) is the sum of the individual pressures (collisions of only one kind of particle) of each gas present VAPOUR PRESSURE OF WATER AT VARIOUS TEMPERATURES

- during a laboratory experiment, it is very common to collect gases by the displacement of water - hydrogen gas is commonly collected by the bubbling of gas (downward displacement of water) into a container filled with water - the gas collected in the container is not pure hydrogen, but a mixture of hydrogen gas and water vapour - water vapour arises from the evaporation of some liquid water - the amount of water vapour is dependent on the temperature of the surroundings - in order to correct for the presence of water vapour, we must use Dalton’s law of  partial pressures

Page 28 of 28

Ex. A volume of 110.0mL of hydrogen is collected over water at a temperature of  17.0 degrees Celsius and an atmospheric pressure of 95.0kPa. what is the volume of the dry hydrogen at STP? P1 = 95.0kPa

STP:

Pwater = 1.94kPa

P2 = 101.325kPa

V1 = 110.0mL

T2 = 273K  

T1 = 17.0 degrees Celsius

V2 = ?

= 290K  Ptotal = 95.0kPa – 1.94kPa = 93.06kPa

V2 =(

)( ) =( = 95.1mL

)(

)