Chemistry AS - ABG Flashcards
What is an element?
A substance that cannot be broken down chemically into anything simpler
What is an atom?
The smallest possible particle of an element
What is a compound?
A substance composed of 2 or more elements chemically bonded together
What is an ion?
A particle formed by an atom or group of atoms losing or gaining electrons
What is a molecule?
An arrangement of atoms covalently bonded together
What is a covalent bond?
A shared pair of electrons holding atoms together in a molecule
What is an ionic bond?
Electrostatic attraction between oppositely charged ions
What is the relative atomic mass of an element?
The average mass of an atom of an element on a scale where an atom of carbon-12 12
What is the relative isotopic mass of a substance?
The mass or an atom of an isotope of an element of a scale where an atom of carbon-12 is 12.
What is the relative molecular/formula mass of a substance?
The average mass of a molecule or formula unit on a scale where an atom of carbon-12 is 12
What is a Mole
An amount of a substance that contains exactly the same number of particles as there are atoms in 12g of C-12
What is stoichiometry?
The relative number of moles of each substance that react together
What is the molecular formula of a substance?
Shows the number of of atoms of each element in one molecule of a substance
What is the empirical formula of a substance?
Shows the simplest whole number ratio of atoms of elements present in a substance
What is a salt?
A compound formed when the H+ ion from an acid is replaced by a metal or ammonium ion
What are acids?
Proton donors - produce H+ ions in aqueous solution
What are bases?
Proton acceptors - remove H+ ions In Aqueous solution
Name the 2 main indicators used in titrations, and give observations
1) Methyl orange - turns yellow to red when acid is added to an alkali 2) Phenolphthalein - red to colourless when acid is added to alkali
What is spin pairing?
Electrons in orbitals spinning in different direction
What shape is a s-orbital?
Spherical
What shape is a p-orbital?
Dumbbell
How many p orbitals are there in one sub-shell?
3, 2 electrons in each, orbitals at right angles to each other
What is first ionisation energy?
The energy needed to remove one electron from one mole of gaseous atoms to form one mole of gaseous 1+ ions
How does nuclear charge affect ionisation energy?
More protons in nucleus -> more positively charged nucleus -> stronger attraction for electrons -> higher IE
How does distance from nucleus affect ionisation energy?
Attraction falls off rapidly with distance -> e- close to nucleus -> much more strongly attracted -> higher IE
How does shielding affect ionisation energy
As the number of electrons between outer electrons and nucleus increases, shielding increases, so attraction between outer and electrons decreases -> IE is lower
What is second ionisation energy?
Energy needed to remove 1 electron from each ion in 1 mole of gaseous 1+ ions to form 1 mole of gaseous 2+ ions
Why do successive ionisation energies increase within each shell?
Electrons are being removed from an INCREASINGLY POSITIVE ION - LESS REPULSION amongst remaining electrons, so the are HELD MORE STRONGLY by the nucleus
When do big jumps in successive ionisation energies increase, and why?
When a new shell is broken into - electron is being removed from a shell closer to the nucleus, and therefore is more strongly attracted to the nucleus
Why do ionic compounds only conduct electricity when molten or dissolved?
Ions in a liquid are free to move and carry a charge When solid they are held in place by strong ionic bonds
Why do ionic compounds have high melting points?
The giant ionic lattices are held together by strong electrostatic forces Lots of energy is needed to overcome these forces and break the ionic bonds
Why are ionic compounds soluble in water?
Water is polar, and attracts the ions away from the ion, forming ion-dipole complexes
What shape do the atoms in a diamond crystal lattice (covalent) have?
Tetrahedral
Which properties of diamond can be attributed to its strong covalent bonds?
Very high melting point - sublimes Hard Good thermal conductor - vibrations travel easily Won’t conduct electricity - all outer electrons are held in localised bonds Won’t dissolve in any solvent
What are allotropes?
Different forms of the same element in the same state e.g. Diamond and Graphite
Describe the structure of graphite
Each carbon has 3 covalent bonds, forming a sheet of flat hexagons each carbons’s 4th electron is delocalised between the hexagon sheets the sheets are bonded together by weak Van Der Waal’s forces
How does graphites structure give rise to its different chemical properties?
1) weak bonds between layers of graphite are easily broken - slippery, used as a dry lubricant, used in pencils 2) delocalised electrons are free to move & conduct an electric current 3) Far apart layers - less dense than diamond - makes strong, lightweight sports equipment 4) Strong covalent bonds - very high melting point - insoluble
describe giant metallic lattice structures:
1) delocalised electrons in metal atoms outer shell - electrons free to move - positive metal ion left 2) +ve metal ions attracted to -ve metal ions - form lattice of +ve ions in sea of -ve electrons - metallic bonding
How does metallic bonding explain metals properties?
1) Number of delocalised electrons per atom affects mpt: - more d.e, stronger bonding, higher mpt - (size of metal ion and lattice structure also affect mpt) 2) no bonds holding ions together, metal ions can slide past eachother - metals are malleable and ductile 3) delocalised electrons pass kinetic energy to each other - good thermal conductors 4) delocalised electrons carry a current - good electrical conductors 5) Strong metallic bonds - insoluble except liquid metals
How does electronegativity make a bond polar?
bonding electrons are pulled towards the more electronegative atom the greater the difference in electronegativity, the more polar the bond
what are the 3 types of intermolecular forces?
1) Hydrogen bonding (strongest) 2) Permanent dipole-dipole interactions 3) Van der Waals forces
Why does reactivity increase down group 2?
As you go down the group: 1) Ionisation Energies decrease - Increasing atomic radius and shielding 2) when grp 2 elements react they lose electrons - lower I.E. = easier to lose electrons = more reactive
What is produced when group 2 metals react with water? give an equation
Metal hydroxide formed M (s) + 2(H2O) (l) -> M(OH)2 (aq) + H2 (g)
What is produced when group 2 metal Oxides react in water?
Aqueous metal hydroxide ions MO (s) + H2O (l) -> (M)2+ + 2OH- (aq)
why do group 2 Oxides form more alkaline solutions as you go down the group?
the hydroxides become more soluble
What is formed when group 2 elements burn in oxygen?
solid white oxides form 2M + O2 -> 2MO
what is thermal decomposition?
when a substance breaks down when heated
give an equation for the thermal decomposition of a group 2 metal
MCO3 -> MO (s) + CO2
how does thermal stability change down group 2?
thermal stability increases - more heat is required to break it down
give examples of uses for group 2 metals
1) Ca(OH)2 - slaked lime - used in agriculture to neutralise acid soils 2) Mg(OH)2 -antiacid - indigestion
give the equation for neutralisation
H+ + OH- -> H2O
what colour and state is fluorine at rtp?
yellow gas
what colour and state is chlorine at rtp?
green gas
what colour and state is bromine at rtp?
red-brown liquid
what colour and state is Iodine at rtp?
grey solid
Describe and explain the pattern of change in mpt/bpt as you go down group 7?
1) Boiling/ melting points increase down group 7 2) size and relative mass of atoms increase: = strength of VDW’s increases
describe and explain the change in reactivity as you go down group 7
Less reactive as you go down group 7 1) halogens react by gaining an electron - reduced -> oxidising agents 2) as you go down group 7, atom size increases -> outer electrons further from nucleus 3) outer electrons are shielded from increasingly positive nucleus - harder for larger halogen atoms to attract an electron 4) larger atoms = less reactive = less oxidising
What is the general ionic equation for reacting Silver Nitrate with a Halogen?
Ag+ (aq) + X- (aq) -> AgX (s)
describe the method for testing for halides:
1) Add Nitric acid to halide solution - removes unwanted ions 2) add AgNO3 (aq) 3) a coloured precipitate forms 4) add NH3
describe the observations when the following halide ions undergo the test: 1) Cl- 2) Br- 3) I-
1) Cl- : White ppt, dissolves in dilute NH3 2) Br- : Cream ppt, dissolves in conc. NH3 3) I - : yellow ppt, insoluble in NH3
What is Sodium Chlorate (I) (NaClO) used for?
Bleach - water treatment, bleaching paper and textiles
What is the equation for the formation of bleach? what conditions are needed?
2NaOH (aq) + Cl2 (aq) -> NaClO (aq) + H2) (l) rtp
What is the equation for the disproportionation of chlorine in water
Cl2 (g) + H2O (l) <–> HCl (aq) + HCLO ( Chloric (i) acid)
Give the equation of the ionisation of Chloric (I) acid:
HClO (aq) + H2O (l) <–> ClO- (aq) + H3O+
Why is chlorine used in water treatment?
1) Kills disease-causing microorganisms 2) prevents reinfection in water supply 3) prevents growth of algae - eliminating bad smells and tastes - removes discolouration
What are the disadvantages of adding chlorine
1) Chlorine gas is very harmful - irritates respiratory system 2) Liquid chlorine causes severe chemical burns 3) Chlorine reacting with organic compounds makes Chlorinated Hydrocarbons - carcinogens
Why is fluoride ions added to water in some parts of the uk, and why are some people concerned by it?
1) Fluoride ions prevent tooth decay 2) small amount of evidence linking fluoride to bone cancer 3) toothpaste is already fluorinated
On what occasions can Hydrogen Bonding occur, and Why?
1) When H is covalently bonded to F, N or O 2) H has a high charge density (v. small atom) , others are very electronegative 3) Bond is polarised 4) H of one molecule can form a weak bond with F,N or O on another molecule
How does hydrogen bonding affect the physical properties of substances, compared with non polar molecules of similar sizes?
Hydrogen bonding makes molecules: 1) Soluble in water 2) Higher Boiling and Melting points
How are molecules of H2O held together in ice?
by H bonds, in a lattice
Why is ice less dense than water?
H bonds are relatively long
