unit 1: periodicity and properties of elements Flashcards
(128 cards)
1
Q
what is an atom?
A
- the smallest particle of a chemical element that can exist
2
Q
what is an element?
A
- a substance consisting of atoms which all have the same number of protons i.e. the same atomic number
3
Q
what is meant by atomic number?
A
- the number of protons in the nucleus of an atom, which characteristic of a chemical element and determines its place in the periodic table
4
Q
what are isotopes?
A
- atoms of the same element with the same number of protons, but different number of neutrons and hence different mass numbers
5
Q
what is meant by mass number?
A
- the total number of protons and neutrons in the nucleus of an atom
- different isotopes of the same element have different mass numbers
6
Q
what is meant by relative atomic mass?
A
- the mean mass of the atoms of an element compared with 1/12 of the mass of a carbon-12 atom
- it is an average of the mass number of all the different isotopes of that element
7
Q
how is the relative atomic mass (Ar) calculated?
A
- Ar = ( isotopic mass x % abundance ) +
( isotopic mass x % abundance ) / 100 - multiply the mass number of each isotope by it’s relative abundance
- add them all together
- divide by 100 if abundance is in %
8
Q
why is the relative atomic mass not always a whole number?
A
- different isotopes of the same element have different mass numbers and the relative atomic mass is an average of the mass numbers of all these isotopes
9
Q
what was the Bohr theory?
A
- an atom has a small, positively charged central nucleus, orbited by electrons at fixed energy levels i.e. distances from the nucleus
10
Q
describe the structure of an atom in the nuclear model of an atom.
A
- in the nuclear model of an atom it displays a small, positively charged central nucleus a long with neutrons that are orbited by electrons
11
Q
what is one difference between the Bohr model and the nuclear model?
A
- in the nuclear model, electrons orbit the nucleus without specific distances
- while in the Bohr model, electrons move in fixed orbits at distinct energy levels
12
Q
what is the structure of an atom?
A
- a small, positively charged central nucleus which contains protons and neutrons, orbited by electrons in shells
13
Q
what is in the nucleus of an atom?
A
- the positively charged central core of an atom, containing protons and neutrons and nearly all of its mass
14
Q
what are the relative masses and charges of protons, neutrons and electrons?
A
proton = 1, +1
neutron = 1, 0
electron = 0.0005 or 1/1836, -1
15
Q
how can the number protons, neutrons and electrons of an element be calculated?
A
protons = atomic number
neutrons = difference between mass number ( big ) and atomic number ( small )
electrons = number of protons
16
Q
how are elements arranged in the periodic table?
A
- in order of increasing atomic number, in rows called periods ( for s and p block, period number = number of electron
shells ) - elements with similar properties are placed in the same vertical columns called groups ( for s and p block elements, group number = number of outer shell electrons )
17
Q
what ions and charges do each group on the periodic table form?
A
- group 1 = +1
- group 2 = +2
- group 3 = +3
- group 4 = can be either negative or positive due to them being transition metals
- group 5 = -3
- group 6 = -2
- group 7 = -1
18
Q
what is a sub-shell?
A
- electrons are found in energy levels
( shells ) around the nucleus, however each shell is made of sub shells and electron configurations reflect these
19
Q
what are the different types of sub-shells?
A
s sub-shell - groups 1 and 2 including helium and hydrogen
d sub-shell - groups 3-12
p sub-shell - groups 13-18
20
Q
how many electrons can each sub-shell hold?
A
s sub-shell can hold 2 electrons
p sub-shell can hold 6 electrons
d sub-shell can hold 10 electrons
21
Q
what is the order for filling sub-shells and give an example?
A
order = 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p
example:
- carbon, 1s², 2s², 2p² as it contains 6 electrons
22
Q
give examples of a solid.
A
- sodium chloride
- sodium carbonate
- sodium
- calcium
- magnesium
23
Q
give an example of a liquid.
A
- bromine
- water
24
Q
give an example of a gas.
A
- oxygen
- carbon dioxide
- steam
25
give an example of an aqueous solution.
- nitric acid
- sulphuric acid
- hydrochloric acid
- magnesium chloride solution
- potassium nitrate solution
- sodium carbonate solution
26
what is a compound?
- a substance that is composed of two or more seperate elements
27
what is a molecule?
- a group of atoms bonded together, representing the smallest fundamental unit of a chemical compound that can take part in a chemical reaction
28
what is an ion?
- an atom or molecule with a net electric charge due to the loss of gain of one or more electrons
29
how is a positive ion formed?
- a positive ion is formed when a neutral atom loses one or more electrons from it's valence shell
- positive ions are name 'cations'
30
how is a negative ion formed?
- a negative ion is formed when a neutral atom gains one or more electrons from it's valence shell
- negative ions are named 'anions'
31
what is a monoatomic ion?
- a charged particle constituting of one atom
- formed by the gain or loss of electrons to the valence shell ( the outer most electron shell ) in a single atom
32
what are some examples of ions and what are there formulas?
OH- = hydroxide
CO₃²⁻ = carbonate
NO3- = nitrate
SO4²- = sulphate
PO4³- = phosphate
NH4+ = ammonium
33
what is a molecular ion?
- formed by the gaining or losing of elemental ions such as a proton, H+
- the four main ions are carbonate, sulphate, ammonium and nitrate
34
what are group 7 elements called when they become an ion?
fluorine = fluoride
chlorine = chloride
bromine = bromide
iodine = iodide
35
what is relative formula mass?
- the sum of the relative atomic masses of an element e.g. water contains 2 hydrogen and 1 oxygen so Mr would be 1 + 1 + 16 = 18
36
what is a mole?
- the amount of any substance containing ( 6.02 x 10 ^23 ) particles
37
what is molar mass?
- the mass per mol of a substance in g mol-1
38
what formula is used to calculate the number of moles of an element?
- number of moles (mol) = mass (g) / molar mass ( g mol-1 )
- n = m / Mr
- e.g. how many mols are there in 117g of NaCl
= 117 / (23 + 25.5 = 58.5 ) = 2 mol
39
what's the formula for calculating molar mass and relative formula mass?
- m = n x Mr
- Mr = m / n
40
what is empirical formula?
- shows the simplest whole number ration of atoms of each element in a compound
e.g. P4O6 can be simplified to P203 but C12H22O11 cannot be simplified as it cannot be divided by the same number when simplifying, so will stay the same
41
what is meant by molecular formula?
- shows the number and type of atoms present of each element in a compound
42
how do we work out empirical formula?
- use the mass given in the questions
- find the number of moles using
n = m / Mr
- divide the number of moles by the smallest value to get a ratio
- adjust the ratio to make them whole numbers, for example if the ratio was 1:1:1.5, double everything so it becomes 2:2:3
- then add the numbers from the ratio to make a formula
43
example of empirical formula
2.00g of a compound contains 1.20g of magnesium and 0.80g of of oxygen. what is the empirical formula?
- mg = 1.2g / 24.3 = 0.049
- O = 0.8g / 16 = 0.05
- 0.049 / 0.049 = 1
- 0.05 / 0.049 = 1.02 rounded down is 1
- ratio 1:1
- empirical formula is MgO
44
what happens if the mass is presented as a percentage?
- if the mass is presented as a percentage e.g. 75% carbon and 25% hydrogen, the mass would be 75g and 25g
- if the question only displays one percentage, e.g. 45% carbon, then do 100-45= 55 to get 55% hydrogen
45
what is meant by the term stoichiometry?
- the ratio, of the amount in moles, of each substance in a chemical reaction
46
what equation do we use when calculating stoichiometry?
n = m / Mr
moles = mass / molecular formula of the element
47
how do we calculate stoichiometry if the question is asking for mass?
1. make sure the equation is balanced
2. then figure out the moles of the given element / compound using the mass in the question
3. once the moles of the given is calculated, put the given element and unknown element into a ratio e.g. 1:1 or 2:1
4. if the ratio is 1:1 the moles of the given and unknown are the same, if the ratio is 2:1 times by 2, if the ratio is 3:2 then do the mass in the question / (2x3)
5. then use m=n x Mr to calculate the mass of unknown
48
give an example of how to calculate stoichiometry?
2NaClO3 -> 2NaCl + 3O2
how many grams of NaCl are produced when 80 grams of O2 are produced?
1) n = 80 / ( 16 x 2 ) = 2.5 mol of O2
2) 2.5 / 3x2 ( ratio of oxygen to NaCl ) = 1.67
mol of NaCl
3) mass = Mr x n = ( 23 + 35.5 ) x 1.67 = 97.70g NaCl
49
give another example when the moles are given instead of mass
2H2 + O2 -> 2H2O
if 3.00 moles of H2O are produced, how many grams of oxygen must be consumed?
1) 2:1 ( 2H2O : O2 ) so we half 3.00 moles to get 1.5 mol of O2
2) Mr of O2 = 16 x 2 = 32 gmol⁻¹
3) mass = 1.5 x 32= 48g of O2
50
what is an electron shell?
- a group of atomic orbitals with the same principal quantum number, n, also known as a main energy level
51
what is a sub-shell?
- a group of orbitals of the same type within a shell
52
what is each sub-shell made up of?
- made up of one type of atomic orbital only ( e.g. s, p, d or f )
53
what is a orbital?
- a region within an atom that can hold up to two electrons with opposite spins
- they have different types
( s, d and f ) and different shapes ( spherical, dumb bells etc. ) and we use box diagrams to represent electrons in orbitals
54
how may electrons can the 1st, 2nd and 3rd shell hold?
1st shell- 2 electrons
2nd shell- 8 electrons
3rd shell- 18 electrons
55
what does the number of electron shells determine?
- what PERIOD the element is in
56
what does the number of outer electrons determine?
- what GROUP the element is in
57
how do we fill shells and sub-shells?
- orbitals are represented by boxes
- electrons are represented by arrows
1. electrons are added one shell at a time
2. the lowest energy level
( 1s ) is filled first
3. each energy level must be filled before the next level
4. each orbital is filled singly before pairing ( fill each box with one arrow first then double them up depending on how many electrons are left over )
58
give an example of how to write the electron configuration for an atom
example: fluorine
- fluorine contains 19 protons / electrons
- to write an electronic configuration you must follow: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d and so on
e.g. fluorine would have an electronic configuration of 1s², 2s², 2p5
59
what is meant by the first ionisation energy?
- the minimum energy required to remove an electron from the valence shell of an isolated gaseous atom to form a 1+ ion
- this energy is measured as a gas
60
what is ionisation?
- when an atom loses an electron from it's outer shell
61
write an equation, including state symbols, to show the reaction that occurs in the first ionisation energy of lithium, aluminium and calcium
1) Li (g) -> Li+ (g) + e-
2) Al (g) -> Al+ (g) + e-
3) Ca (g) -> Ca (g) + e-
62
what does the size of ionisation depend on?
- the atomic radius ( size of the atom )
- this changes depending on where the element is in the periodic table
63
what is the atomic radius?
- this is the size of the atom
- atomic radius decreases across a period
- atomic radius increases down a group
64
what happens to the first ionisation energy across a period?
- it increases across a period
65
what happens to the nuclear charge and the atomic radius across a period?
- the nuclear charge increases across a period - each successive element has one more proton than the last
- therefore, attraction of electrons to the nucleus increases
- atomic radius decreases
( so electrons are closer to the nucleus )
66
what happens to the first ionisation energy down a group?
- it decreases down a group
- there are more inner shell electrons as you go down the group, so there is shielding of outer electrons
- atomic radius increases
( so electrons are further away )
- therefore, attraction of electrons to the nucleus decreases
67
what 3 things does atomic radius depend on?
1) electron shells- the more shells, the larger it's atomic radius
2) nuclear charge- more protons in the nucleus means more attraction between electrons and the nucleus so atomic radius decreases ( electron closer to nucleus )
3) shielding- atomic radius increases with the amount of electrons as more shielding from electrons means less attraction between the nucleus and electrons
68
what is an ionic bond?
- ionic bonds are strong electrostatic attractions between positive and negative ions
- an ionic bond is between a metal and non-metal
69
what is the structure of an ionic compound?
- ionic compounds have giant structures
- in a giant ionic structure, the ions are arranged in a regular, three-dimensional pattern called a lattice
- the electrostatic forces between the ions act in all directions and keep the structure together
- the large number of these strong electrostatic attractions gives ionic compounds high melting points
70
give an example of a strong ionic compound.
- sodium chloride lattice
- in this lattice, each sodium ion is surrounded by 6 chloride ions and each chloride ion is surrounded by 6 sodium ions
- this repeating pattern continues for a vast number of ions
71
how are positive ions formed?
- positive ions are generally formed by metal atoms losing electrons
- they will have a positive charge equal to the group number if formed from a group 1,2,3 element
- they will have a different charge if formed from a transition metal e.g. Fe ²⁺
- are also known as cations
72
how are negative ions formed?
- negative ions are generally formed by non-metal atom gaining electrons from metal ions
- they will have a negative charge equal to 8 minus the group number of the elements
- are also known as anions
73
what is the atoms radius?
- refers to the mean or typical distance from the centre of an atoms nucleus to the outermost isolated electron
74
how do we convert between degrees and kelvins?
degrees -> kelvin = +273
kelvin -> degrees = -273
75
what is a covalent bond?
- it is an electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms
76
how do atoms form covalent bonds?
- a covalent bond forms when atoms share a pair of electrons
- the electron pair is attracted to both nuclei and is localised between them
- generally each atom in the bond contributes one electron to the pair, but a covalent bond consisting of an electron pair derived from one of the atoms is called a dative covalent
( coordinate ) bond
77
give an example of a giant covalent bond.
- graphite
- each carbon atom is bonded to three other carbon atoms and is around 3600 degrees so has a high melting point
- it is a good conductor of electricity as is has a structure that allows electrons to move freely ( delocalised electrons )
- insoluble in water
78
how do we work out how many covalent bonds an element will form?
e.g. oxygen
- take the group number of oxygen, which is group 6 and take it away from 8
8-6 = 2 covalent bonds formed
79
why do noble gases not form covalent bonds?
- as they already have a full outer shell and therefore do not want to form covalent bonds with any other element
80
how are bond strength and bond length related?
- they are inversely related
- this meaning that the shorter the covalent bond length, the greater the covalent bond strength e.g. a double bond with 2 electrons is stronger than a single bond
81
what are the strengths of covalent bonds?
- they have a giant covalent lattice structure meaning they have high melting and boiling points as many strong covalent bonds need to be broken to melt them e.g. diamond which is a lattice of carbon
82
what is an intermolecular force?
- intermolecular forces are interactions between molecules caused by either permanent or induced dipoles
83
what are the 3 types of intermolecular forces?
1. London dispersion forces
(temporary dipole-induced dipole)
2. Dipole-Dipole forces
(permanent-dipole-dipole bonds)
3. Hydrogen bonding
84
order the intermolecular bonds from weakest to strongest.
weakest- temporary-dipole-induced dipole
middle- permanent-dipole-induced dipole
strongest- hydrogen bonding
85
how does a temporary dipole form?
- e.g. in 2 xenon atoms their electrons are constantly moving
- at any one instant the electron cloud may be distributed unequally, this causes partial charges δ+ and δ– to develop and a temporary dipole forms
86
what happens at a temporary dipole?
- the temporary dipole on one atom induces a dipole on a neighbouring atom
- there is an electrostatic force of attraction between the δ– atom and the δ+ on the other atom
- these are temporary dipole-induced dipole attractions, they are present between all molecules but they are weak
87
what happens to the boiling point as you go down a group?
- there are more electrons in an element as you go down the group
- this means there are stronger London dispersion forces
- this means more energy is required to seperate the intermolecular forces
- boiling point then increases
88
what happens as you increase the chain length of an alkane?
- the boiling point increases
89
what is electronegativity?
- the ability of an atom to attract bonding electrons in a covalent bond
- permanent dipole-dipole attractions arise due to electronegativity
90
what happens to electronegativity as you go down a group?
- it decreases
- e.g. H = 2.1, Li = 1.0, Na = 0.9
91
what happens to electronegativity as you go across a period?
- it increases
- e.g. Li = 1.0, Be = 1.5, B = 2.0, C = 2.5 etc.
92
how can molecules be represented?
- in 3D using wedges and dashes
- straight lines show bonds in the same plane as the paper
- wedges indicate a bond coming out of the plane of the paper
- dashes indicate a bond going into the plane of the paper
93
what makes a molecule polar?
- if a molecule has polar bonds it cannot be symmetrical in every plane
- therefore is has a permanent dipole and is a polar molecule
94
what makes a molecule non-polar?
- if a molecule has polar bonds but they are arranged in opposing directions then they cancel out and the molecule is non-polar
95
what is hydrogen bonding?
- this is an intermolecular force between an electron-deficient hydrogen atom δ+ and a lone pair on oxygen, nitrogen or fluorine atoms
- O, N, and F are the only atoms that can form hydrogen bonds as they are small and highly electronegative, which means they pull pairs of electrons towards them
- water molecules can form hydrogen bonds between each other
- hydrogen bonds are strong intermolecular forces but still much weaker than a covalent bond
96
give an example of how a hydrogen bond forms between 2 hydrogen atoms and an oxygen atom.
- a water molecule contains 2 hydrogen atoms and an oxygen atom
- each hydrogen atom is bonded to the oxygen atom by a single covalent bond
- oxygen is more electronegative than hydrogen, so it gains a partial negative charge
- the 2 hydrogen atoms gain a partial positive charge
- there are permanent dipole-permanent dipole attractions between water molecules
- they occur in a particualr set of circumstances in water, and are called hydrogen bonds
- oxygen has 6 electrons in it's outer shell, two of these form bonding pairs with the electrons from the 2 hydrogen atoms, leaving 4 electrons that are not involved in bonding so form lone pairs of electrons
- there is a force of attraction between a lone pair of electrons on an oxygen atom and the δ+ charge on a hydrogen atom so forms a hydrogen bond
97
what other compounds have hydrogen bonding?
- ammonia, water and hydrogen fluoride
- if the molecule has London dipole forces and a hydrogen bond the boiling point will be higher than just a molecule with London dipole forces
98
what is the difference between ice and liquid water?
- when ice freezes the water forms a giant lattice structure
- it is less dense than liquid water which is why it floats
99
why does water have a high boiling and melting point?
- as it has extensive hydrogen bonding between water molecules as well as large London dipole forces, therefore resulting in larger intermolecular forces and more energy to overcome the forces of attraction
100
what is a metallic bond?
- metallic bonds occur between metal atoms
- this kind of bonding is present in any metal element or alloys
( mixtures of different metals and other substances )
- the metallic bond is a strong electrostatic attraction between the positive metal ions and the delocalised electrons
101
what is the structure of a metallic bond?
- the positive ions are layered in 3 dimensions, so form a giant lattice structure
102
give an example of a giant metallic structure.
- magnesium
- contains ions and is a good conductor of electricity as it has free moving electrons
- has a moderate melting point of around 650 degrees
- not soluble in water
103
what are the properties of metals?
1) high thermal and electrical conductivity due to the delocalised electrons, which are free to move
2) high boiling and melting point due to strong electrostatic attractions between positive ions and electrons
3) malleability, can be shaped, as layers of positive ions slide over each other and the delocalised electrons move with the layers so strong metallic bonds remain intact
4) ductility, can be pulled into wires as positive ions roll over each other and the delocalised electrons move with the positive ions, so strong metallic bonds remain intact
104
what are delocalised electrons?
- these are the electrons from the outer shell of the metal atoms, but are not fixed to a particular atom, so can move freely throughout the structure
105
what is the pattern of melting points as you go down group 1?
- the melting points decrease as the atoms gets larger, as larger metals have more electrons and more electron shells meaning they have more shielding between the nucleus and delocalised electrons
- this means that the electrostatic force between them is weakened producing a weaker metallic bond, so less energy is required to break these forces
106
why do group 2 metals have a higher melting point than group 1 metals?
- each group 2 metal has a higher melting point than the group 1 metal in the same period
- this is because the group 2 metal has 2 delocalised electrons per positive ion, rather than one so this gives it a greater electron density around the positive ions
- this, alongside with +2 charge, produces a stronger electrostatic attraction between the nucleus and the delocalised electrons, and so a stronger metallic bond
107
what is an oxidation number?
- a measure of the number of electrons that an atom uses to bond with an atom of another element
- oxidation numbers are derived from a set of rules also known as an elements oxidation state
108
what are the first 4 rules of oxidation numbers?
rule 1- the oxidation number of a neutral element is zero e.g. H2, F2, Na, and O2
rule 2- the oxidation number of a monatomic ( one-atom ) ion is the same as the charge on the ion e.g. Na+ has an oxidation number of =1, Cl- has an oxidation number of -1 and S2- has an oxidation number of -2
rule 3- the sum of the oxidation numbers in a neutral compound is zero and the sum of all oxidation numbers in a polyatomic ( many-atom ) ions is equal to the charge on the ion e.g. the sum of the oxidation numbers in Na2CO3 is 0
rule 4- in compounds the elements of:
- group 1 has an oxidation number of +1, G2 = +2, G3 = +3 etc.
109
what are the next 5 rules of an oxidation number?
rule 5- the oxidation state of hydrogen in a compound is usually +1 e.g. HCl, H2SO4 all have oxidation numbers of +1
rule 6- fluorine always has an oxidation number of -1 in compounds e.g. NaF, SnF have oxidation numbers of -1
ruler 7- the oxidation number of oxygen in a compound is usually -2 e.g. MgO
rule 8- in transition metals the oxidation number can vary e.g. in Fe2O3 the oxidation number is +3 but in FeO it is +2
rule 9- chlorine, bromine and iodine usually have oxidation numbers of -1, except when it is in a compound with oxygen e.g. HCl= -1 but NaClO3 is +5
110
what is a redox reaction?
- a reaction that involves reduction and oxidation
O- oxidation
I- is
L- loss of electrons
R- reduction
I- is
G- gain of electrons
111
what happens to the oxidation numbers during oxidation and reduction?
oxidation- it loses electrons so the oxidation number increases
reduction- it gains electrons so the oxidation number decreases
112
write a half equation for the oxidation of magnesium.
Mg -> Mg²⁺ + 2e-
- magnesium loses 2 electrons
113
write a half equation for the reduction of oxygen.
O2 + 2e- -> O²-
- oxygen gains 2 electrons
114
what is the equation to work out the concentration in an aqueous solution?
- concentration is a measure of the amount of solute dissolved per unit of solvent
concentration (mol dm-3) = amount (mol) / volume (dm3)
c = n / v
115
how do we calculate percentage yield?
percentage yield= actual amount (mol) of product / theoretical amount (mol) of product x100
116
what do period 2 and 3 elements form when they react with oxygen?
- oxides often when being burnt in pure oxygen
e.g. Li2O, Na2O
117
what is the general trend in melting and boiling points across a period and why?
-melting and boiling points increase across groups 1 to 4 and decrease across groups 5 to 0 as the type of bonding changes from metallic to covalent across the period
-the melting and boiling points of metals increase as the metallic bond strength increases because the positive metal ions get smaller across the period, their charge increases and there are more delocalised electrons, so there is a stronger attraction between the cations and the delocalised electrons
-the elements from group 4 ( carbon and silicon) both have very high melting and boiling points as they have giant covalent structures with many strong covalent bonds
-the elements from groups 5 to 0 are simple molecules / atoms with only weak London forces between them
118
what happens in a metal displacement reaction?
- a more reactive metal displaces a less reactive metal from a metal salt
- the more reactive metal atoms gain electrons to form ions, so are reduced
- each metal ion loses electrons to form atoms, so is oxidised
119
what happens in a halogen displacement reaction?
-a more reactive halogen ( further up group
7) displaces a less reactive halogen from a halide salt
-the more reactive halogens gain one electron per atom to form halide ions, so are reduced
-each halide ion loses an electron to form halogens, so is oxidised
120
when do the oxides tend to form alkaline solutions and why?
-if the metal oxides dissolve in water as they form metal hydroxides
-e.g. sodium oxide reacts to form sodium hydroxide
121
what are the reactions of metals with water?
-some metals (e.g. those in groups 1 and 2) react directly with water to form hydroxides
-the metals in group 1 react immediately on contact with water, fizzing as hydrogen is produced, dissolving to form an alkaline solution, and sometimes catching on fire or exploding
-the metals in group 2 react in a similar but slower way, e.g. magnesium needs to be reacted with steam to form hydrogen
-less reactive metals don't tend to react directly with water, but may react with water and oxygen to corrode
122
what are the reactions of metals with dilate acids?
-metals that are more reactive than hydrogen ( potassium, sodium, lithium, calcium, magnesium, aluminium, zinc and iron) will react with dilute acids to form a salt and hydrogen gas
-the salt forms when a metal ionises and metal ions replace hydrogen ions in the acid
-the salt formed depends on the dilute acid
used (sulfuric acid = sulfate, hydrochloric
acid = chloride )
123
what are the trends in metals reactivity and why?
-as you go down a group, the metals become more reactive with oxygen, water and dilute acids because the atoms become larger in size
-as you go across a period, the metals become less reactive with oxygen, water and dilute acids because more electrons have to be lost to form metal ions
124
what is a transition metal?
-a d-block element which forms at least one stable ion with an incomplete d-subshell
125
what are some examples of uses of nonmetal oxides?
-boron trioxide ( B203) = glass manufacture,
e.g. optical fibres
-carbon dioxide ( CO2 ) = to make drinks fizzy
and in fire extinguishers to prevent oxygen gas reaching flames
-nitrous oxide ( N20 ) = pain relief, e.g
during childbirth
-silicon dioxide ( SiO2 ) = food additive to
stop powders sticking together
-phosphorus pentoxide ( P205 ) = removes
water from organic molecules in the chemical industry
-sulfur dioxide ( SO2 ) = manufacture of
sulfuric acid ( contact process )
126
what are some examples and uses of metal oxides and hydroxides?
-aluminium oxide (Al203) = in abrasive
paper as its giant ionic lattice structure makes it very hard
-magnesium hydroxide ( Mg(OH)2 ) = antacid
medicine as it neutralises excess HCl in the stomach
-sodium hydroxide (NaOH) = drain cleaner
as it will react and breakdown fats and oils from food waste
127
what are some examples and uses of metal salts?
-sodium chloride (NaCl) = food industry as
flavouring and a preservative, and raw material in chemical industry to make hydrogen, chlorine and sodium hydroxide through electrolysis as it's readily avaliable and very soluble in water
-potassium sulfate ( K2SO4) = fertiliser as
very soluble and contains minerals needed by plants
-magnesium sulfate ( MgSO4 ) and sodium
sulfate (Na2SO4 ) = a drying agent as their
giant ionic crystalline structure can absorb water molecules
-calcium sulfate ( CaSO4 ) = key component
of plaster used to cover walls in buildings
128
what are some examples of uses of transition metals?
-copper = in electrical wiring as its metallic
structure allows electrons to move freely and so conduct electricity effectively
-vanadium = added to steel as vanadium
steels are very hard and resistant to wearing so can be used in engines
-titanium = in aircraft manufacture as it's as
strong as steel but much less dense
