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1

what is matter and the three different states

- matter: anything that occupies space and has a mass, three states of matter
- solid: fixed volume / shape, no compressibility, particles vibrate about fixed positions, strong cohesive forces, very little thermal expansion, high density
- liquid: fixed volume, variable shape, almost no compressibility, particles move past one another, high density, limited thermal expansion
- gas: variable volume / shape, very compressible, particles move freely, low density, little to no cohesive forces, high thermal expansion

2

what is an element, compound, mixture and molecule

- element: pure, homo atomic molecules or individual atoms, cannot be subdivided, fixed properties
- compound: pure, hetero atomic molecules, chemically subdivided (elements / simpler compound), fixed properties, chemically combined in fixed ratio
- mixture: combination of pure substances, elements or compounds, hetero (variable) / homo (uniform) geneous, properties vary, easily separated (filtration, evaporation)
- molecule: smallest particle of pure substance, capable of independent existence

3

what are physical vs chemical properties

- physical properties: observed or measured without changing composition (colour, density, melting point)
- physical changes: changing physical dimensions of matter without changing composition (cutting, crushing, dissolving, heating)
- chemical properties: properties expressed when changed to different compound (splitting compounds, burning), some are reversible others not

4

what is daltons atomic theory

- matter is made up of atoms
- all atoms of an element are identical and different from atoms of other elements
- compounds are combinations of atoms of two or more elements
- molecules of a compound contain the same number of atoms of each element found in the compound
- atoms are rearranged in chemical reactions (separated / combined), never created or destroyed

5

describe an atom, atomic number, atomic mass and isotopes

- atom: smallest particles of matter, made up of particles / sub-atomic particles (arrangement makes atoms different for different elements), nucleus and 1 or more electrons
- atomic number: protons or electrons, equal, ions occur when these aren’t equal, above element
- atomic mass: protons + neutrons in nucleus, protons decide what element it is (electrons can change)
- isotope: same atomic number (protons), different mass number (different neutrons)
- relative atomic mass: takes abundance of isotopes and their masses into account (on periodic table)

6

describe elements of the periodic table / periodic law

- structure: vertical columns (groups) and horizontal rows (periods)
- elements: arranged in terms of increasing atomic numbers
- periodic law: elements in this order ensure similar chemical properties occur at regular / periodic intervals
- metals: high melting point, hard, ductile, malleable, shiny, conduct electricity / heat
- nonmetals: solids, liquids or gases, brittle, soft or hard, dull, don’t conduct electricity / heat

7

describe electrons, their configuration and valence electrons

- electrons: negatively charged particles, orbit nucleus in shells / orbits, position dependent on energy
- movement: from one shell to another by absorbing (further away) or releasing (original position) energy
- shells / sub-shells: higher shell (higher energy, further from nucleus), vice versa,
- valence electrons: participate in chemical reactions forming bonds, gain / lose electrons for full shell
- noble gases: inert, full valence shell, stable atoms, don’t react (He, Ne, Ar)
- lewis diagram: represents valence electrons of an atom, ionic (transfer) and covalent (sharing)

8

what are the three types of bonding

- ionic: metal (low ionisation) and nonmetal (high ionisation), negative and positive ions bind, leading to transfer of electrons (donator / acceptor) and production of two stable ions (cation / anion)
- covalent: sharing of electrons, two molecules deficient (no donator / acceptor), each donates e- which are jointly shared (single, double or triple bonds), nonmetals
- metallic: force between positive nuclei and negative delocalised valence electrons (loosely bound, readily realised into lattice), metals only, high conductivity

9

what is the difference between binary ionic and binary covalent and ionic

- binary ionic: metal and non metal (two different elements), diatomic (two molecules, may or may not be the same), name cation first (specifying the charge, if necessary), then the nonmetal anion (suffix -ide)
- binary covalent: nonmetals, two elements can combine in several different ways (prefixes necessary), first (neutral name), second (suffix -ide), prefix used in front of both to indicate number of atoms
- prefixes: mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca, if there is only one of the first element in the formula, the mono- prefix is dropped
- ionic: metal and polyatomic atom (ions composed of two or more atoms linked covalently), cation first (specifying the charge, if necessary), then the polyatomic ion (following given names)

10

what is / are the types of chemical reactions

- chemical reaction: transformation of one or more substances into one or more new substances
- balancing equations: number and type of atoms on both sides must be equal
- acid base: acid + base → water + salt
- acid carbonate: acid + carbonate → water + salt + carbon dioxide
- acid metal: acid + metal → hydrogen gas + salt
- precipitation: solution + solution → precipitate
- decomposition of metal carbonate: metal carbonate → metal oxide + carbon dioxide

11

describe the three types of relative masses

- relative atomic (Ar): mass of an atom compared to 1/12 of the mass of carbon 12 (exactly 12), no units
- relative molecular (Mr): mass of molecule compared to 1/12 of the mass of carbon 12, adding Ar of atoms
- relative formula (Mr): mass of formula unit of iconic compound compared to 1/12 of the mass of carbon 12, adding Ar of atoms

12

what is stoichometry and the calculation involved

- stoichiometry: quantitative relationship of chemical reactions between reactants and products
- moles (n): amount of a substance that contains the same number of atoms as there in 12g of C12, mol
- avogadro’s constant (NA): number of atoms in 12g C12 (6.022x1023)
- molar mass (M): mass of one mole of a substance, equal to Ar or Mr mass (m) in grams (g)
- steps: write balance equation, determine molar ratio, convert masses to moles, ratio of unknown : known for new moles, convert moles to masses

13

what are limiting reagents

- limiting reagent: reactant which is used up first, hence ending the chemical reactant, stoichiometric ratio (theoretical, balanced equation) and actual ratio (actual, number of moles present)
- steps: determine moles of reactants, use equation for SR (x / y), use actual number of moles to determine AR (moles x / moles y = x), if equal no limiting reagent if not vice versa


14

what is / are the types of energy

- Ep: distance between particles, greater distance = higher energy, solids < liquids < gases, absorbed energy overcomes attractive forces increases as melting / vaporises occurs, increases with increased space between particles (gas)
- Ek: velocity of particles, temperature increases particle velocity increases, Ek = 1⁄2 mv2, increases as solids, liquids and gases are heated, lowest as the temperature decreases
- energy transfer: heat is a form of energy which is usually transferred in a chemical reaction, exo / endo
- units of energy: metric units are joule (J) and kilojoule (kJ), others include calorie (cal), kilocalorie (kCal)
- calorimetry: measure of how much heat is gained or lost

15

what is a heating / cooling curve

- plot of temperature vs time for a substance when energy is supplied constantly
- phase change: represented by a flat line where there is no change in temperature as the temperature supplied is being used to break the bonds between particles / atoms allowing the state change to occur

16

describe evaporation, vapour pressure and boiling

- evaporation: change from liquid to gas at a temperature below boiling point, surface of liquid
- vapour pressure: pressure exerted by gas molecules as they escape surface of liquid / solid, different liquids have different vapour pressures at the same temperature (difference in cohesive forces)
- boiling: change from liquid to gas at a substances boiling point, occurs when vapour pressure = atmospheric pressure, occurs throughout whole liquid sample

17

what is a solution, solute and solvent

- solution: substance is dissolved in a second substance to form a homogeneous mixture / solution
- solute: substance that is dissolved (smaller quantity), continue to exist as molecules surrounded by solvent
- salt: dissolve in water into positive and negative ions (dissociation)
- solvent: substance that does the dissolving (often water)

18

what is solubility and the three qualitative states

- solubility: qualitative, maximum amount of solute dissolved in a specific amount of solvent under given temp and pressure
- unsaturated: less than maximum amount of dissolved solute at given temp and pressure
- saturated: contains maximum amount of dissolved solute at given temp and pressure
- supersaturated: unstable, tat contains more than saturated, saturating at higher temp then cooling it

19

what are the different ways solution concentration can be calculated (5)

- concentration: quantitative, amount of solute in a solvent
- moles per litre: mol L-1) = n (mol) / V (L)
- grams per litre: g L-1 = solute (g) / solution (L)
- % composition by mass: % composition = (mass of solute / mass of solution) x 100
- % composition by volume: % composition = (volume of solute / volume of solution) x 100
- parts per million (ppm): C (ppm) = mass of solute (mg) / mass of solution (kg)

20

what are types of energy reactions

- exothermic: release heat to surroundings, enthalpy of products less than reactants, ΔH -ve, hot, heat is a product
- endothermic: absorb heat from surroundings, enthalpy of products more than reactants, ΔH +ve, cold, heat is a reactant
- spontaneous: occurs at room temp without addition of continuous energy, generally exo (lower energy)
- state changes: solid to liquid to gas (endothermic) and gas to liquid to solid (exothermic)

21

what is reaction rate and important features of it

- reaction rate: rate at which reactants are used up, Δ [reactant] / time or Δ [product] / time
- collision theory: reactants must regularly collide with sufficient energy and favourable orientation to allow for bonds to be broken and new ones formed
- activation energy (Ea): minimum amount of energy that must be absorbed to cause a reaction
- transition state: activated complex, unstable, bonds of reactants are broken and bonds of products formed
- ions (faster reaction) and solids (slower reaction)

22

what are factors that affect rate of reaction

- nature of reactants: if no bonds need to be broken = rapid (precipitation) vice versa
- reactant concentration: increased conc, increases RoR due to increased collisions and hence increased successful collisions with favourable orientation
- subdivision: increased SA, increases RoR due to increased area where collisions can occur
- temperature: increased temp increases RoR due to increased velocity of molecules and increased collisions, increased collisions with sufficient energy to overcome Ea barrier, hence successful collisions
- catalyst: lowers Ea, increasing number of molecules with sufficient energy to overcome barrier (enzymes)

23

what is chemical equilibrium

- equilibrium: reversible chemical reaction that has constant macroscopic properties and rate of forward reaction = to reverse
- equilibrium constant (Keq): relationship between [reactants] and [products] at equilibrium
- Keq = [products] / [reactants], only gases / aqueous solutions are used in the expression
- changes to system in equilibrium will only effect substances in the Keq
- magnitude: extent of reaction, high Keq (products > reactants) and low Eq (reactants > products)

24

what is le chateliers principle and what are factors that affect it

- LC principle: if a chemical system at equilibrium is subject to a change in conditions, the system will adjust to re-establish equilibrium and partially counteract imposed change
- concentration: promotes reaction that will counteract change (same as changes in partial pressure)
- volume / pressure: decreasing volume increases concentration / gas pressure, system favours reaction with least number of gaseous particles and vice versa
- temperature: increased temperature favours endothermic and decreased temperature favours exothermic
- catalyst: no effect, increases both forward and reverse reaction, increases rate that equilibrium is reached

25

what are the two acid base theories

- arrhenius: acids produce hydrogen ions in solution and bases produce hydroxide ions
- bronsted lowry: when reacting an acid and a base, a conjugate acid will be produced from the base (proton acceptor) and a conjugate base produce from the acid (proton donator)

26

describe features of acids

- properties: turns blue litmus red, sour tasting, conduct electricity in solution (presence of ions), ions act as charge carriers (proton donator)
- ionisation constant (Ka): extent of ionisation, equilibrium constant, HA ↔ H+ + A-, Ka = [H+][A-] / [HA]
- strong: completely ionise in aqueous solution, large Ka, HCl, H2SO4, HNO3, HBr
- weak: partially ionise in aqueous solution, small Ka, CH3COOH, HF, NH4, H2CO3, H3PO4
- reactions: with metals, carbonates / hydrogen carbonates and metal oxide

27

describe features of bases

- properties: turns red litmus blue, tastes bitter, feels soap / slippery to touch, conduct electricity in solution, ions act as charge carriers (proton acceptor)
- strong: completely dissociate in aqueous solution, all group 1-2 metal oxides / hydroxides, are Kb
- weak: partially dissociate, small Kb, NH3, CO32-
- reactions: with acids (neutralisation), amphoteric metals (Zn, Cr, Al to produce H2 and salt containing complex ion) and amphoteric hydroxides (salt containing complex ion)

28

what is the self ionisation of water

- water: weak electrolyte, 2H2O (l) ↔ H3O+ (aq) + OH- (aq)
- Kw: [H+][OH-] = (1.0x10-7) (1.0x10-7) = 1.0x10-14 at 25C, can be used to calculate [H] and [OH]

29

describe pH and pH calculations

- pH: describe the acidity / alkalinity of solution, < 7 (acidic), 7 (neutral) and > 7 (basic), pH + pOH = 14
- calculation (pH): pH = -log10[H+] where [H+] = concentration mol L-1
- calculation (pOH): pOH = -log10[OH-] where [OH-] = concentration mol L-1
- finding [H+]: [H+] = 10^-pH (pH 7 = 10^-7 [H+])

30

what are buffers

- buffer: solution that can resist changes in pH when small amount of acids or bases are added, consist of equimolar amounts of weak acid and its conjugate base
- add acid: system favours the reaction that effectively removes the added H3O and maintaining pH
- add base: system favours the reaction that effectively removes the added OH and maintaining pH
- capacity: extent to which buffer can absorb added acid or base, determined by [buffer ions]