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foundations in chemistry Flashcards

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1
Q

atomic number

A

the number of protons in the nucleus

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2
Q

mass number

A

total number of protons and neutrons in the atom

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3
Q

number of neutrons =

A

mass number - atomic number

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4
Q

isotopes

A

atoms of the same element with the same number of protons but different numbers of neutrons

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5
Q

isotopes have similar…
isotopes have varying…

A

similar chemical properties as they have the same electronic structure
varying physical properties as they have different masses

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6
Q

relative isotopic mass

A

the mass of one isotope compared to one-twelfth of the mass of one atom of carbon-12

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7
Q

relative atomic mass

A

the weighted mean mass of one atom compared to one-twelfth of the mass of one atom of carbon-12

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8
Q

relative molecular mass

A

the average mass of a molecule compared to one-twelfth of the mass of one atom of carbon-12

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9
Q

how to calculate relative atomic mass

A

= total(isotopic mass x % abundance)/100

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10
Q

ammonium formula

A

HN4+

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11
Q

phosphate formula

A

PO4 3-

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12
Q

carbonate formula

A

CO3 2-

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13
Q

sulfate formula

A

SO4 2-

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14
Q

nitrate formula

A

NO3-

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15
Q

hydroxide formula

A

OH-

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16
Q

what are spectator ions

A
  • ions that are not changing state
  • ions that are not changing oxidation number
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17
Q

ions definition

A

have a different number of electrons and protons

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18
Q

mole definition

A

the amount of substance in grams that has the same number of particles as there are atoms in 12 grams of carbon-12

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19
Q

molar mass

A

the mass in grams of 1 mole of a substance
given the unit g/mol-1

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20
Q

how to calculate molar mass

A

add up all the mass numbers of each element in the compound

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21
Q

molar gas volume definition

A

units = dm3mol-1
the volume of 1 mole of a gas at a given temperature and pressure.
at 1atm and 25degrees the molar gas volume is 24dm3mol-1

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22
Q

avogadro’s constant definition

A

there are 6.02x10 23 atoms in 12 grams of carbon-12

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23
Q

what can avogadro’s constant be used for

A

atoms, molecules and ions

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24
Q

mole =

A

mass/mr

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25
Q

empirical formula definition

A

the simplest ratio of atoms of each element in the compound

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26
Q

molecular formula definition

A

the actual number of atoms of each element in the compound

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27
Q

concentration =

A

n/v (moles/volume)
conc = moldm-3
volume = dm3

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28
Q

cm3 to dm3
cm3 to m3
dm3 to m3
m3 to dm3

A

divide by 1000
divide by 1000000
divide by 1000
times by 1000

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29
Q

formula for a new diluted concentration

A

new diluted concentration = original concentration x (original volume/new diluted volume)

30
Q

ideal gas equation and units

A

PV = nRT
pressure = Pa
Volume = m3
Temp = K
n=moles
R = 8.31JK-1mol-1

31
Q

how to convert from celsius to kelvin

A

add 273

32
Q

molar gas volume equation

A

v(dm3) = n x 24

33
Q

number of particles/ions/molecules/atoms equation

A

moles x avogadro’s constant

34
Q

density equation
and density units

A

density = mass/volume
density units = gcm-3

35
Q

percentage yield equation

A

actual/theoretical x 100

36
Q

how can %yield be lowered

A
  • incomplete reactions
  • side reactions
  • losses during transfer of substances
  • losses during purification stages
37
Q

atom economy equation

A

mass of useful products/mass of all reactants x 100
USE BALANCING NUMBERS

38
Q

acid definition

A

releases H+ ions in aqueous solutions
proton (H+) donor

39
Q

common strong acids
common weak acid

A

HCL
H2SO4
HNO3
CH2COOH (ethanoic acid) - weak acid

40
Q

base definition

A

proton (H+) acceptor

41
Q

common bases

A

metal oxides
metal hydroxides
ammonia

42
Q

alkali definition

A

soluble base that releases OH- ions in aqueous solutions

43
Q

common alkalis

A

sodium hydroxide (NaOH)
potassium hydroxide (KOH)
ammonia (aq) - NH3

44
Q

strong and weak acid definitions

A

strong = completely dissociate when dissolved in water
weak = slightly dissociate when dissolved in water, to form an equilibrium mixture

45
Q

salt definition

A

formed when the H+ ion of an acid is replaced by a metal ion or an ammonium ion

46
Q

acid + base =
acid + carbonate =

A

salt + water
salt + water + carbon dioxide

47
Q

observations in a carbonate reaction

A
  • effervescence due to CO2 gas
  • solid carbonate will dissolve
48
Q

oxidation

A

the process of electron loss
involves an increase in oxidation number

49
Q

reduction

A

process of electron gain
involves a decrease in oxidation number

50
Q

oxidation numbers for…
group 1 metals
group 2 metals
Al
H
F
Cl, Br,I
O

A

+1
+2
+3
+1 (except in metal hydrides where it is -1 eg NaH)
-1
-1 (except in compounds with oxygen and fluorine)
-2 (except in peroxides where it is -1 and compounds with fluorine)

51
Q

electron structure

A

1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4p6, 5s2, 4d10…

52
Q

shapes of s sub levels and p sub levels

A

s sub levels = spherical
p sub levels = dumbbells

53
Q

ionic bonding definition

A

electrostatic force of attraction between oppositely charged ions formed by electron transfer

54
Q

what makes ionic bonding stronger and the melting points higher?

A
  • when the ions are smaller and have higher charges
55
Q

what are the physical properties of ionic compounds?
give an example of the structure

A
  • high melting points and bp: strong electrostatic forces between oppositely charged ions
  • non conductor of electricity when solid: ions are held in lattice and so cannot move
  • conductor of electricity when molten or in solution: ions are free to move and charge can be carried
  • soluble in aq solutions
    FORMS GIANT IONIC LATTICES
56
Q

covalent bonding definition

A

the strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms

57
Q

dative covalent bonding definition

A

forms when the shared pair of electrons in the covalent bond comes from only one of the bonding atoms

58
Q

covalent bonding properties and types of structure formed

A

covalent bonding forms SIMPLE MOLECULAR structures
low mp and bp: due to weak intermolecular forces between molecules
solubility: poor
conductivity: poor as no ions to conduct and electrons are fixed in place

59
Q

electronegativity definition

A

the relative tendency of an atom in a covalent bond in a molecule to attract electrons in a covalent bond to itself

60
Q

most electronegative atoms are…

A

F,O,N and Cl

61
Q

factors affecting electronegativity…

A
  • increases across a period as number of protons increases and the atomic radius decreases because the electrons in the same shell are pulled in more
  • decreases down a group as the distance between nucleus and outer electrons increases the shielding of inner shell electrons increases
62
Q

what does a small electronegativity difference indicate?

A

the compound is covalent

63
Q

how does a permanent dipole (polar covalent) bond form

A

when the elements in the bond have different electronegativities

64
Q

whats a polar covalent bond

A
  • has an unequal distribution of electrons in the bond, producing a charge separation (dipole)
65
Q

where do induced dipole-dipole interactions occur

A
  • between all simple covalent molecules
  • between the separate atoms in noble gases
  • NOT IN IONIC SUBSTANCES
66
Q

how do induced dipole-dipole interactions form

A
  • in a molecule, the electrons are constantly moving
  • as this happens, electron density fluctuates and parts of the molecule become more or less negative
  • from this, temporary dipoles form
  • these temporary dipoles cause dipoles to form in neighbouring molecules. these are induced dipoles.
67
Q

factors affecting the size of induced dipole-dipole interactions

A
  • the more electrons, the higher the chance that temporary dipoles will form. this makes the induced dipole-dipole interactions stronger and so boiling points will increase.
68
Q

explain permanent dipole dipole forces

A
  • occurs between polar molecules
    -they are stronger than induced dipole–dipole interactions and so the compounds have higher boiling points
  • Polar molecules have a permanent dipole. (commonly compounds with C-Cl, C-F, C-Br H-Cl, C=O bonds)
  • Polar molecules are asymmetrical and have a bond where there is a significant difference in electronegativity between the atoms.
69
Q

where does hydrogen bonding occur

A
  • in compounds that have a hydrogen atom attached to one of the three most electronegative atoms of nitrogen, oxygen and fluorine, which must have an available lone pair of electrons
  • it occurs in addition to induced dipole dipole interactions
70
Q

explain the molecular structure of ice

A

Water can form two hydrogen
bonds per molecule, because
oxygen is very electronegative,
and it has two lone pairs of
electrons

The molecules are held
further apart than in liquid
water and this explains the
lower density of ice

71
Q

explain the molecular structure of iodine

A
  • covalent bonds between the iodine atoms in the I2 molecule
  • crystals contain a regular arrangement of I2 molecules held together by weak induced dipole-dipole interactions

chem retake (2 decks)

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