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Secondary School Fall Semester 2018 Written For Tumblr 11TH GRADE Chemistry ___ Notes Introduction: History of the Atom Average Atomic Mass 👌 Mass represented on the periodic table calculated based on the abundance of each isotope (weighted average) ○ general formula: average mass =( %a x mass a) + (% b x mass b)... ■ ** always express the % as a decimal, 58% as 0.58) Lewis Dot Diagrams and Structures Lewis suggested a means of keeping track of outer electrons which are called valence electrons There are 2 bond types : ionic and covalent 👌 In ionic bonding one atom has a stronger attraction for electrons than the other, and “steals” an electron from a second atom ○ Ionic bonds form most often between nonmetals and metals 👌 In covalent bonding the attraction for electrons is similar for two atoms. They share their electrons to obtain an octet. ○ If two atoms have approximately the same pull on electrons, they share the electrons (forming a “covalent” bond) ○ Non-metals form covalent bonds with each other ■ CO32- and NH4+ are both composed of covalent bonds even though they are often found as part of an ionic compound, there are a couple of tricks to figure out how to show the covalent bonds of polyatomic ions Electronegativity 👌 Not all atoms share electrons equally in a covalent bond. Some have a greater attractiveness for electrons which results in the electrons being shared spending more time near the more electronegative atom. 👌 To determine how equally electrons are shared in a covalent bond we find the difference in electronegativity between the atoms that form the bond. ○ If the difference is between 0.5 and 1.7 we refer to the bond as polar. Molecular Shape and Molecular Polarity Recall a polar bond occurs when the atoms involved in the bond have an E.N between 0.5 and 1.7 This results in a polar bond where one end is partially positive the other is partially negative. In simple diatomic molecules having a polar bond results in a polar molecule, this is not true for more complex molecules in these cases we must consider the geometry of the molecule Intramolecular and Intermolecular Forces Intramolecular forces : occur within a molecule or compound Intramolecular Force Basis of Formation Relative strength Ionic Cations to anions 1 Polar covalent Partial charge cation to 2 partial charge anion Non-polar covalent Nuclei shared electrons 3 Intermolecular forces : attractive forces between molecules or between molecules and ions 👌 Weaker than intramolecular forces but still important because they determine the physical properties of molecules like their boiling point, melting point, density, and enthalpies of fusion and vaporization ○ Dipole-Dipole Forces (DDF) ■ Occurs only between polar molecules ■ occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule ■ Dipole-dipole interactions are the strongest intermolecular force of attraction ○ Hydrogen Bonding ■ Occurs only between oxygens and hydrogens or nitrogen or fluorine ■ Hydrogen bonding is a relatively strong force of attraction between molecules, and considerable energy is required to break hydrogen bonds. ■ This explains the exceptionally high boiling points and melting points of compounds like water ○ London Dispersion Forces (LDF) ■ Occurs between both polar molecules and non-polar molecules ■ These force occur when the electrons in an atom happen to be located on one side of the molecule. ■ This is a temporary effect ■ The more electrons a molecule has, the stronger the London dispersion forces Properties of Ionic Compounds 👌 Contain ion - ion attraction as well as LDF 👌 High melting point and boiling point - you must overcome the attractive forces between the ions 👌 Soluble - ionic compounds dissociate in water….break apart into their component ions 👌 Electrical conductivity - only conductive as a aqueous solution Properties of Covalent Compounds containing hydrogen bonds 👌 These compounds contain hydrogen bonding, as well as LDF, their melting and boiling points are lower than ionic 👌 Soluble in water due to the attraction between the dipoles on the water and the other polar compound. Properties of Polar Covalent Compounds 👌 These compounds contain dipole-dipole interactions as well as LDF, their melting and boiling points are lower than molecules contain hydrogen bonds 👌 Solubility in water due to the attraction between the dipoles on the water and the other polar compound. ○ Larger polar covalent molecules are less soluble as the attraction between the same molecules is stronger. Properties of Non-polar Covalent Compounds 👌 These compounds contain LDF, their melting and boiling points are lower than polar covalent molecules 👌 Not soluble in water as there is limited attraction between the water and the non-polar covalent molecule Properties of molecular compounds 👌 Electrical conductivity- cannot conduct because of absences of ions 👌 Exception - acids undergo a processes call ionization when dissolved in water and are therefore weakly conductive. Chemical Nomenclature 👌 SIMPLE BINARY COMPOUNDS ○ contain 2 elements ○ first element has a positive valence; second has a negative valence ○ name of second ends in –ide ○ first element has only 1 common valence ■ barium oxide 👌 VARIABLE VALENCE BINARY COMPOUNDS ○ These compounds have a first element with 2 (or more) possible positive valences (Refer to Table 1). Three systems have been developed to resolve this problem. ■ IUPAC METHOD ● a Roman numeral is used to indicate the positive valence ○ tin (IV) chloride ■ –ous / -ic METHOD’ ● names are often based on the Latin names (cuprum, ferrus, stannum, plumbum, aurum, mercurium) ● the –ous suffix indicates the lower valence ● the –ic suffix indicates the higher valence ○ ferric oxide, ferrous oxide ■ Greek Prefix Method ( can be used for metals but we will only use it with Molecular Compounds only) ● Greek prefixes can be used to indicate the number of atoms in a compound (unless there is only one of the first element) ● do not need to consider valences of the elements ☺ ● can be used for any binary compound (usually nonmetals) ○ nitrogen trifluoride, silicon dioxide 👌 BINARY ACIDS ○ acids are compounds that release hydrogen ions (H+) ○ binary acids contain 2 elements; first is hydrogen ○ dissolved in water (aq = aqueous) *always include (aq) with chemical formula for an acid ■ Name: hydro_________________ic acid ● hydrofluoric acid, hydrosulfuric acid 👌 BASES ○ contain the polyatomic ion hydroxide (OH-) ○ a polyatomic ion is a group of atoms with a net charge ○ these ions are treated as if they consist of 1 element ■ sodium hydroxide, lead (IV) hydroxide 👌 Elements ○ Noble Gases are always monoatomic (1 atom in formula) ○ other gases / halogens are diatomic (2 atoms bonded) ○ metals are always monoatomic in their elemental form ■ helium gas, nitrogen gas 👌 SALTS OF POLYATOMIC IONS ○ The term “salt” refers to any ionic crystalline compound. One common group of salts are known as binary compounds ○ Polyatomic ions are simply collections of atoms that are covalently bonded together but carry a net charge, such as hydroxide ([OH]-) ○ Since these groups of atoms form ions, they can also form salts with oppositely charged ions ■ sodium nitrate, ammonium sulfide ○ Polyatomic Ions that contain oxygen can have a variety of different formulas because the number of oxygen atoms can vary ■ Hypo___ite (remove 2 0) ■ ________ite (remove 1 0) ■ ________ate ■ Per_____ate (add 1 O) ○ However exceptions occur ■ Sulfate, phosphate 👌 OXYACIDS ○ OXYACIDS are formed between hydrogen ions and polyatomic ions containing oxygen. ○ To avoid confusion with binary acids, oxyacids have their own nomenclature system. ■ hypo___ous acid (remove 2 0) ■ ________ous acid (remove 1 0) ■ ________ic acid ■ per_____ic acid (add 1 O) ● hypoiodous acid, bromic acid, sulfurous acid 👌 ACID SALTS ○ these are composed of part acid and part salt ○ (metal) + (hydrogen + polyatomic ion) ■ barium hydrogen carbonate, copper (II) dihydrogen phosphite 👌 PEROXIDES ○ contain the ion [O2]2- ○ these compounds have one more oxygen than a normal binary oxide ○ do not reduce these formulas ■ calcium oxide vs calcium peroxide 👌 HYDRATES ○ salts with water molecules loosely bonded ○ number of water molecules given by Greek prefix ■ Calcium sulfate dihydrate, Copper (II) sulfate pentahydrate Chemical Reactions 👌 Synthesis reactions ○ occur when two substances (elements) combine and form a compound. ■ Two elements forming a new binary compound ■ a univalent metal reacts with a non-metal to form an ionic compound ● when predicting the product it is important to refer to the ions formed during the electron transfer not the subscripts on the molecular reactants ■ a multivalent metal reacting with a non-metal to form various compounds ● It is important to consider all possible products of the reaction using any possible charges of the multivalent metal ■ Two compounds forming a new compound ● Non-metal oxide reacting with water forms an oxy-acid ● these types of reactions are directly responsible for acid precipitation ■ Water and a metal oxide always forms a base 👌 Decomposition Reactions ○ occur when a compound breaks up into the elements or in a few to simpler compounds ■ a binary compound decomposes into its elements ● electrolysis, which is a process where electrical energy is used to propel a chemical reaction ● decomposition does not often happen naturally heat can also be used ■ a metal nitrate decomposes into a metal nitrite and oxygen gas ■ a metal carbonate decomposing into a metal oxide and carbon dioxide ● the decomposition of calcium carbonate into calcium oxide is a major step in the curing of cement ■ a metal hydroxide decomposing into a metal oxide and water ● occurs typically when heat is applied 👌 Combustion Reactions ○ In general: complete combustion ■ CxHy + O2 → CO2 + H2O + energy ○ or combustion of a metal ■ M (S) + O2(g) + energy → MO(s) 👌 Single Displacement Reactions ○ occur when one element replaces another in a compound. ○ A metal can replace a metal (+) OR a nonmetal can replace a nonmetal (-) ■ element + compound→ product + product A + BC → AC + B (if A is a metal) OR A + BC → BA + C (if A is a nonmetal) ■ (remember the cation always goes first!) ○ how do we know if the reaction will happen or not? ■ The Activity Series: The higher on the list the the metal is the more likely it will replace the metal atom in the compound ○ Non-Metal displacing another non-metal ■ Halogens are the most common example of this type of single displacement reaction 👌 Double Displacement Reactions ○ occur when a metal replaces a metal in a compound and a nonmetal replaces a nonmetal in a compound ■ Compound + compound → product + product ■ AB + CD → AD + CB ○ a reaction that forms a solid ■ many double displacement reactions result in the formation of a precipitate, an insoluble product ○ a reaction that forms a gas ■ Some double displacements result in the formation of a gas, often H2, O2, NH3 and CO2 ■ Often the formation of the gas is often a second step in the reaction process ● CH3COOH(aq) + NaHCO3(s) → NaCH3COO(aq) + H2CO3(aq) decomp of carbonic acid H​2​CO​3​(aq) → CO​2​(g) + H​2​O(l) ○ we typically write this reaction as one step CH3COOH(aq) + NaHCO​3​(s) → NaCH​3​COO(aq) + CO​2​(g) + H​2​O(l) *Notice that this doesn’t really look like the generic double displacement reaction ■ Reactants Products acid + compound containing ionic compound + water+ carbon carbonate ion dioxide compound containing ammonium ionic compound + water + ions + ammonia compound containing hydroxide ions ○ a reaction that forms water ■ These double displacement reaction are difficult to detect because while a new product is formed the product is water which makes it difficult to detect ■ This type of reaction typically takes place between an acid and a base and is called a neutralization reaction ● acid + base → water + salt reactant 1 + reactant 2 → products include acid + has carbonate ions → water and carbon dioxide aqueous salt + aqueous salt → a precipitate has ammonium ions + has hydroxide ions → water and ammonia acid + has hydroxide ions → water The Mole 👌 A counting unit 👌 Similar to a dozen, except instead of 12, it’s 602 billion trillion 602,000,000,000,000,000,000,000 👌 6.02 X 10​ (in scientific notation) 23​ A Mole of Particles: 6.02 x 1023 ​ ​ particles 👌 n = ​N N​A ○ n = amount in moles (mol) N = the number of particles N​A​ = avogadro’s number Molar mass: the mass in grams of 1 mole of an element or compound written in grams per mole 👌 The masses of the elements on the periodic table are all determined based on the fact that one atom of C - 12 has a mass of 12 amu, this means that 1 mole of C-12 has a mass of exactly 12g Percent Composition: Based on the Law of Definite Proportion, a compound always has the same percentage composition 👌 % composition =​mass of an element mass of a compound x 100 Chemical Formulas of Compounds: give the relative numbers of atoms or moles of each element in a formula unit - always a whole number ratio (the law of definite proportions) 👌 Empirical Formula: The formula of a compound that expresses the smallest whole number ratio of the atoms present ○ Ionic formula are always empirical formula 👌 Molecular Formula: The formula that states the actual number of each kind of atom found in one molecule of the compound Stoichiometry Calculating Quantities in Chemical Reactions Table template: Molecular Ratio Molar Mass Mole Ratio Mass (g) of reactants and products Limiting and Excess Reactants Excess Reactant: The reactant that remains at the end of the reaction. Percentage Yield: A reaction rarely yields the amount of product we theoretically predict, theoretical yield. The actual yield, that is determined through experiment is impacted by reaction conditions such as temperature and pressure as well as purity and quality of the reactants. ● % yield = ​actual yield​ x 100% theoretical yield SOLUTIONS & SOLUBILITIES Solution: a homogeneous mixture containing particles the size of a typical ion or covalent molecule. (0.1–2.0 nm in diameter) Aqueous solution: a solution that contains water Solute: the dissolved substance in a solution Solvent: the major component in a solution Solubility: the maximum amount of solute that will dissolve in a given quantity of solvent at a specific temperature 👌 A solution is s​ aturated​ when no additional solute can be dissolved at a particular temperature 👌 AS ​ upersaturated​ solution can form when more than the equilibrium amount of solute is dissolved at an elevated temperature, and then the supersaturated solution is slowly cooled. 👌 An ​Unsaturated​ solution is formed when more of the solute can dissolve in it at a particular temperature 👌 When two liquids are completely soluble in each other they are said to be ​Miscible 👌 The amount of solute per unit solvent required to form a saturated solution is called the solute's ​Solubility 👌 The formation of a solution is dependent on the following factors ○ Forces that attract particles of the solute to each other ○ Forces that attract particles of the solute to the solvent ○ Forces that attract particles to the solvent to each other 👌 For solutions in water hydrogen bonding is a key intermolecular force to consider 👌 Solubility of an ionic compound ○ Most are soluble ○ Exceptions occur when the attraction between the charge ions are significantly greater than the attraction to the dipole of the water molecule 👌 Solubility of a molecular compound ○ Polar compounds tend to have high solubility ○ Smaller non-polar compounds have some solubility 👌 Solubility of gases ○ “Like dissolves like” again, polar gases tend to have greater solubility in water in comparison to non-polar gases Solubility is affected by temperature 👌 With increase in temperature solubility of most solids increases 👌 Most gases become less soluble in water as the temperature increases Pressure has little effect on the solubility of liquids and solids. The solubility of gases is strongly influenced by pressure. Gases dissolve more at high pressure Calculating and Expressing Concentration concentration: the ratio of the quantity of solute to the quantity of solvent or the quantity of solution concentrated: having a high ratio of solute to solution dilute: having a low ratio of solute to solution There are many different ways to express concentration including: concentration as a % very small concentrations (ppm, ppb) molar concentration 👌 Concentration as a % ○ mass/volume % (m/v%) ■ describes mass of solute in grams in the total volume of the solution in millilitre ■ m/v % = ​mass of solute (g)​ x 100% volume of solution (ml) ○ mass/mass % (m/m%) ■ describes mass of solute in the total mass of the solution. The mass of the solute and solution must have the same units. ■ m/m % = ​mass of solute ​ x 100% mass of solution ○ volume/volume % (v/v%) ■ describes volume of solute in the total volume of the solution. The volume of the solute and solution must have the same units. ■ v/v % = ​volume of solute ​ x 100% volume of solution 👌 Very small concentrations ○ parts per million (ppm) ■ ppm = ​mass of solute​ x 10​6 mass of solution ○ parts per billion (ppb) ■ ppb =​ mass of solute​ x 10​9 mass of solution ○ The mass of the solute and solution must have the same unit. 👌 Molarity- a chemist's unit of concentration ○ Molarity tells us the strength of a solution in terms of moles/litre mol/L or M it has the symbol C ○ We use molarity to describe solutions – a 1 M solution of NaCl has 1 mole of NaCl per Litre of solution? ■ molarity = ​moles (of solute) ​ C =​ n volume of solution(Litres) V ● volume must always be expressed in litres 👌 Dilution ○ Sometimes we need to decrease the molarity of the sample we are working with, this is called a dilution ■ C​1​V​1​= C​2​V​2 Qualitative Analysis Observation: a statement that is based on what you see, hear, taste, touch and smell 👌 Empirical knowledge – comes directly from observation e.g. water boils at 100​ C o​ Inference: a judgment or opinion that is based on observations and/or conclusions from testing 👌 Theoretical knowledge – comes from our explanation of our observations e.g. we can use the kinetic molecular theory to explain why water boils at 100oC 👌 We can use qualitative analysis to identify unknown elements, ions or compounds. We use knowledge about elements, ions or compounds and our observations to draw conclusions ○ Flame Test ■ If we heat metal ions we can cause them to produce specific colour of flame. (e.g Cu2+ produces a blue-green flame) ■ Fireworks use different metal ions in the form of powder to create the colours ■ You can heat them in solid or aqueous form. ○ Solution Colour ■ Many ionic compounds when in solution will be coloured ■ The colour we observe can help to identify the metal ions and anions in solution ○ Precipitation Reactions ■ We can use solubilities to further confirm the presences of particular ions in solution ■ We can take advantage of insolubility of certain ions to cause precipitate to form 👌 Individually each of the above test is not sufficient to accurately identify an unknown ion(s) in a solution. Used in combination however we can easily determine the presence of specific ions in solution Arrhenius Acids and Bases 👌 The Arrhenius Theory ○ An acid is a substance that ionizes in water to produce one or more hydrogen ions, H+(aq) ○ A base is a substance that dissociates in water to form one or more hydroxide ions, OH-(aq) ■ Acids undergo ionization because they are molecular compounds that form ions when they dissolve in water ■ Bases, which are most often ionic compounds break apart or dissociate into their component ions when dissolved in water ■ The properties of Acids and Bases according to Arrhenius are due to the presence of the hydrogen ion and the hydroxide ion ● So what about a base that does not contain a hydroxide ion? ○ Ammonia is molecular - it ionizes in water ○ NaHCO3(s) when dissolved water becomes H​2​CO​3(aq) + ​ OH​-​(aq) 👌 Properties of Acids ○ They are weak conductors of electricity ○ They turn blue litmus paper: red, bromothymol blue: yellow and phenolphthalein: clear ○ They taste sour, they acts as preservative in foods ○ In strong concentrations they cause burns and blisters on skin ○ When dissolved in water they ionize 👌 Properties of Bases ○ They do not react with metals or carbonates ○ They are good conductors of electricity, ○ They turn red litmus paper blue, bromothymol blue, blue and phenolphthalein pink ○ They taste bitter, ○ In strong concentrations they cause burns and will make your skin feel slippery. ○ When dissolved in water they dissociate 👌 Weak vs Strong ○ Strong acids and bases when dissolved in water dissociate or ionize completely ○ Weak acids and bases when dissolved in water dissociate or ionize only partially* not fully explained by Arrhenius theory ■ The Strong Binary Acids ● all contain a halogen, only HF(aq) does not ionize completely due to the strong electronegativity of the fluorine atom ■ The Strong Oxoacids ● The hydrogens that ionize on these oxoacids are attached to an oxygen, the more oxygens in the oxoacid the stronger the acid. This occurs because oxygen has a high electronegativity. ● diprotic and triprotic acids are those that have more than one hydrogen to lose. ■ Weak Acids ● Common weak acids include acetic acid (CH​3​COOH), hydrocyanic acid HCN, hydrofluoric acid HF and phosphoric acid H​3​PO​4 ■ Strong Bases ● dissociate completely in water to form a metal ion and the hydroxide ion, all alkali metals and alkaline metals below Be all form strong bases ■ Weak Bases ● weak bases interact directly with the water molecules to produce hydroxide ions 👌 Weak vs dilute and strong vs concentrated ○ weak and strong refer to the ease at which hydrogen and hydroxide ions are formed ○ dilute and concentrated refer to the number of hydrogen or hydroxide ions are in the solution- can be described by pH ■ This means you can have a solution that is dilute but contains a strong acid 👌 The pH Scale ○ Used to describe the concentration of hydrogen ions in a solution ○ The higher the concentration of hydrogen ions the lower the pH ○ Therefore the higher the hydrogen ion concentration the more acidic the solution ■ Indicators help you find out whether a solution is acidic or not ● They change colour in acid or base solutions ● Universal indicator changes colour in acids and alkalis ○ pH is a method of report the Hydrogen ion concentration in solution ■ pH = -log [H+] ■ We can calculate the pOH ● Indicates the Hydroxide ion concentration ● a low pOH = a high hydroxide ion concentration ○ Low pOH = a strong base ■ pOH = - log[OH- ] ● pH and pOH are linearly related ○ This means, if we know one, we can easily determine the other ■ pH + pOH = 14 Gases: Properties and Behaviour Changes of State and Forces between particles: Type of Attractive Forces States of Matter Example Between oppositely charged Usually solid Table salt (NaCl) particles Between polar molecules Solid, liquid or gas Glucose *depends on size of Ethanol molecule, the larger the Ammonia molecule the stronger the forces Between non-polar solid , liquid or gas Paraffin molecules *depends on size of Pentane molecule, the larger the Carbon dioxide molecule the stronger the forces 👌 Kinetic Energy of Particles and temperature of a Substance ○ describes the motion of a particle in chemistry we can indirectly measure the kinetic energy of a sample by measuring the temperature ■ The higher the temperature the greater the kinetic energy ○ When we change the state of substance we change the kinetic energy of the substance enough to overcome the intermolecular forces that exist between the molecules 👌 Properties of Gases ○ Gases are compressible ■ The volume of a gas can be decreased significantly by applying pressure. This causes the overall pressure of the gas increases ○ Gases expand as the temperature increases ■ If we increase the temperature of a gas we are increasing the kinetic energy of the sample and if the container can expand it will expand ○ Gases have very low viscosity ■ This means that they move easily from one container to another ○ Gases have much lower densities ■ The density of water vapour is 1/1000 of the density of liquid water ○ Gases are miscible ■ gases will always mix completely with each other unlike liquids(e.g water and oil) 👌 Kinetic Molecular Theory of Gas ○ The theory that explains gas behaviour in terms of random motion of particles with negligible volume and negligible attractive forces ■ Ideal Gas - a hypothetical gas made up of particles that have mass but no volume and no attractive forces between them ○ Gas particles are in constant, random motion ○ Individual gas particles are considered point masses (has no volume) ○ Gas particles have no attractive forces between them ○ Gas particles interact with each other and the walls of their container through only elastic collisions ○ Average kinetic energy is directly related to temperature Boyle’s Law 👌 What is Boyle’s Law? ○ When a gas is under pressure it takes up less space: ■ The higher the pressure, the smaller the volume ○ Boyle’s Law tells us about the relationship between the volume of a gas and its pressure at a constant temperature ○ The law states that pressure is inversely proportional to the volume 👌 Units of Pressure ○ Pressure is force per unit area – in physics it is often represented by N/m​2 ○ In chemistry we represent pressure typically as kPa or atm 1 N/m​2​ = 10​-3 ​KPa = 1.0 x 10​-6 atm ​ ■ 101.325 kPa = 1 atm unit symbol use Equivalent values Standard atm Gas compressors 1.0 atmosphere Millimetres of mmHg Blood pressure 760 mercury torr torr Vacuum pump 760 Pascal Pa Pressure sensor 101 325 bar bar Scuba gear 1.01325 millibar mb barometers 1.01325 x 10​-3 Pounds per square psi tires 14.7 inch 👌 STP and SATP ○ STP stands for standard temperature and pressure = 0o C and 101. 325 kPa or 273 K and 101.325 kPa ○ SATP stands for standard ambient temperature and pressure = 25o C and 100 kPa or 298 K and 100 kPa 👌 Boyle’s Law as a formula ○ Pressure is inversely proportional to the volume and can be written as: Pressure α 1/volume ■ This is more usually written as: Pressure = ​constant volume ■ PV=k ■ P​1​V​1​=P​2​V​2 ● P=pressure in KPa V=volume in L k=constant Gases and Temperature Change 👌 The Kelvin Temperature Scale ○ K = °C + 273. ○ Made because - 273​o​C is known as “absolute zero” ■ At this point all molecular movement stops 👌 Charles’s Law ○ as the temperature of a gas increases, the volume increases proportionally, provided that the pressure and amount of gas remain constant ■ V​1​/T​1​ = V​2​/T​2 👌 Gay-Lussac’s Law ○ the pressure of a fixed amount of gas at constant volume is directly proportional to its Kelvin temperature P ∝ T ■ P/T=k ■ P​1​/T​1​ = P​2​/T​2 ● Temperature should always be in Kelvin 👌 Combined Gas Laws ○ P/T=k or P​1​/T​1​ = P​2​/T​2 ○ V/T = K or V​1​/T​1​ = V​2​/T​2 ○ PV=k or P​1​V​1​=P​2​V​2 👌 Avogadro’s Law ○ a gas law stating that equal volumes of all ideal gases at the same temperature and pressure contain the same number of molecules ■ essentially the volume of gas is directly proportional to the number of molecules of the gas at constant P and V ○ n/V = K ○ n​1​/v​1​ = n​2​/v​2 👌 IDEAL GAS LAW ○ a gas law that describes the relationship among volume, pressure, temperature and moles of an ideal gas ○ universal gas constant: a proportionality constant that relates the pressure on and the volume of an ideal gas to its amount and temperature ○ R = 8.314 KPa L/ mol K ■ PV = nRT ● P = pressure in KPa ● V = volume in L ● n = moles of gas ● R = 8.314 KPa L/ mol K ● T = temperature in K