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ZZ Chem 1 > atomic structure and the periodic table > Flashcards

atomic structure and the periodic table Flashcards

1
Q

Define an atom

A
  • the smallest part of an element that can exist
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2
Q

Define an element

A
  • a substance that only contains one type of atom
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3
Q

Define a compound

A
  • they contain two or more elements chemically combined in fixed proportions
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4
Q

How are compounds formed?

A
  • from elements by chemical reactions
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5
Q

How can compounds be represented?

A
  • by using formula using the symbols of the atoms from which they were formed
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6
Q

What do chemical reactions involve?

A
  • they always involve the formation of one or more new substances
  • and often involve a detectable energy change
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7
Q

How can compounds be separated?

A
  • they can be separated back into their elements by chemical reactions
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8
Q

What is a mixture?

A
  • they consist of two or more elements or compounds that aren’t chemically bonded together
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9
Q

What are the chemical properties of a mixture?

A
  • the chemical properties of each substance in the mixture are unchanged
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10
Q

How can mixtures be separated/

A
  • by using physical separation techniques rather than chemical reactions
  • e.g., filtration, crystallisation etc. * these physical processes do not involve chemical reactions and no new substances are made.
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11
Q

Why are compounds not the same as molecules?

A
  • compounds must be two or more different elements joined together
  • molecules are just two or more of any elements joined together, they could be the same element
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12
Q

What is filtration used to separate?

A
  • it’s used to separate an insoluble solid from a liquid it is suspended (not dissolved) in
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13
Q

Describe the method of filtration

A

1) use a filter funnel and filter paper folded in the shape of a cone
2) pour the mixture into the filter paper and funnel
3) the solid is too large to fit through the pores of the filter paper
4) therefore only the water passes through and the solid remains on the filter paper - so they are separated

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14
Q

What is crystallisation used to separate?

A
  • it is used to obtain and separate a soluble solid from a liquid that it is dissolved in
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15
Q

Describe the method of crystallisation

A
  • leave the solution to evaporate slowly at room temperature
    OR
    1) heat the solution on a bunsen burner, tripod and glaze
    2) this causes the water to evaporate, leaving crystals of the solid behind
    3) ensure that the heat does not affect the solid itself; some solids will undergo thermal decomposition when heated
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16
Q

Why does crystallisation leave behind solid crystals?

A
  • this works because water will have a lower boiling point than the solid dissolved in it,
  • therefore leaving the solid crystals
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17
Q

What is simple distillation used to separate?

A
  • like crystallisation, it separates a soluble solid from a liquid.
  • however, this is used when we actually want to keep the liquid and not simply evaporate it off
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18
Q

Describe the method of simple distillation

A

1) evaporate the liquid by heating it to a high enough temperature using a Bunsen burner
2) condense the vapour by cooling it in another container

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19
Q

What is fractional distillation used to separate?

A
  • it is used to separate a mixture of 2 different liquids
  • these liquids must have different boiling points for this physical separation technique to work
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20
Q

Describe the method of fractional distillation

A

1) connect a flask to a fractionating column full of glass beads and again to a condenser and another beaker
2) heat the flask using a Bunsen burner until the temperature of the thermometer at the top of the column reaches the boiling point of the first liquid
3) at this point, both liquids will be evaporating but only the liquid with the higher boiling point will condense and drop back down once it reaches the top of the column
4) the lower boiling point liquid will not condense and will pass into the condenser, cool and form a pure liquid
5) the remaining liquid will be the other higher-boiling point liquid

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21
Q

What is chromatography used to separate?

A
  • it is used to separate substances in a solution based on their different solubilities
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22
Q

Describe the method of chromatography

A

1) draw a pencil line on the bottom of chromatography paper
2) draw a dot of first colour on pencil line and dot of second colour on the pencil line
3) place the bottom of the paper in a solvent like water and ethanol
4) the solvent makes its way up the paper and the ink is dissolved into it

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23
Q

Why is a pencil line used in chromatography?

A
  • pencil does not dissolve in the solvent
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24
Q

What was the knowledge of the atom before the discovery of the electron?

A
  • atoms were thought to be tiny spheres that could not be divided
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25
Q

What did the discovery of the electron lead to?

A
  • JJ Thomson’s plum pudding model
  • this suggested that the atom is a ball of positive charge with negative electrons embedded in it
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26
Q

What disproved the plum pudding model and how?

A
  • Rutherford’s alpha particle scattering experiment
  • if the plum pudding model was true, the particles should have passed straight through or only be slightly deflected at most because the positive charge was thought to be very spread out throughout the atom
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27
Q

Explain the process of the alpha particle scattering experiment

A
  • positively-charged alpha particles were fired at an extremely thin sheet of gold foil
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28
Q

What were the results of the alpha particle scattering experiment?

A
  • most alpha particles went straight through, showing that the atom is mostly empty space
  • few alpha particles deflected, showing that there is something small and dense in the atom (the nucleus)
  • few alpha particles repelled, showing that there is a small positive charge in the atom (protons); this mass isn’t evenly spread out, meaning the positive charge must be concentrated in the centre
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29
Q

Describe the conclusions of the alpha scattering experiment (which formed the nuclear model)

A
  • the mass of an atom was concentrated at the centre (nucleus)
  • the nucleus was charged (positively), with a cloud of electrons surrounding it
  • the atom is mainly empty space
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30
Q

How did Bohr adapt the nuclear model?

A
  • he adapted the nuclear model by suggesting that electrons orbit the nucleus at specific distances. * the theoretical calculations of Bohr agreed with experimental observations
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31
Q

Why did Bohr conclude that Rutherford’s cloud of electrons wouldn’t work?

A
  • the electrons would be attracted to the nucleus, causing the atom to collapse
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32
Q

After Bohr’s adaptation of the nuclear model, what was discovered?

A
  • later experiments led to the idea that the positive charge of any nucleus could be subdivided into a whole number of smaller particles,
  • each particle having the same amount of positive charge;
  • the name ‘proton’ was given to these particles
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33
Q

What did James Chadwick discover?

A
  • his experimental work provided the evidence to show the existence of neutrons within the nucleus.
  • this was about 20 years after the nucleus became an accepted scientific idea
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34
Q

Number of protons in an atom =

A
  • number of electrons
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35
Q

Define the atomic number

A
  • the number of protons in an atom of an element
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36
Q

Atoms of different elements have different numbers of…

A
  • protons
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37
Q

What is the approximate radius of an atom?

A
  • 0.1nm (1x10⁻¹⁰m)
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38
Q

What is the approximate radius of a nucleus in comparison to atom?

A
  • less than 1/10000 of the atom’s radius (1x10⁻¹⁴m)
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39
Q

Where is almost all of the mass of an atom located?

A
  • in the nucleus
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40
Q

Define mass number

A
  • the sum of the protons and neutrons in an atom
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41
Q

Define an isotope

A
  • an atom of the same element with a different number of neutrons
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42
Q

Describe the nuclear model of the atom

A
  • the nucleus is in the centre of the atom, containing protons and neutrons
  • protons are positively charged and neutrons have no charge (neutral) so the overall charge of the nucleus is positive
  • electrons orbit the nucleus in distinct shells/energy levels
    electrons are negatively charged
43
Q

Define relative atomic mass

A
  • an average mass that takes into account of the abundance of the isotopes of the element
44
Q

What is the relative mass of each subatomic particle?

A
  • proton - 1
  • neutron - 1
  • electron - very small
45
Q

State the RAM formula

A
  • sum of (isotope abundance x isotope mass number) / sum of abundances of all isotopes (100)
46
Q

How do you calculate number of neutrons in the nucleus?

A
  • mass number (number on top) - atomic number (number on bottom)
47
Q

Why does the mass number of the atom not take into account electrons?

A
  • the mass of the atom is concentrated in the nucleus -
  • electrons have negligible mass
48
Q

Describe how the electrons are structured

A
  • the electrons in an atom occupy the lowest available energy levels (innermost available shells).
  • the maximum number of electrons each shell can hold is 2,8,8
49
Q

How can atoms become stable?

A
  • by having a full outer shell
50
Q

How are elements in the periodic table arranged?

A
  • in order of atomic (proton) number
  • elements with similar properties are in columns (groups)
51
Q

What is in between groups 2 and 3 on the periodic table?

A
  • transition metals
52
Q

Why is it called the periodic table?

A
  • because similar properties occur at regular intervals
53
Q

Describe the relationship between elements in the same periodic table group

A
  • they have the same number of electrons in their outer shell,
  • this gives them similar chemical properties and similar reactions
54
Q

What happened before the discovery of protons, neutrons and electrons?

A
  • scientists attempted to classify the elements by arranging them in order of their atomic weights
55
Q

What was wrong with early periodic tables?

A
  • they were incomplete and some elements were placed in inappropriate groups if the strict order of atomic weights was followed (groups where they shared different properties), so these tables were not accepted
56
Q

How did Newlands order his table?

A
  • in order of atomic mass
57
Q

What did Newlands realise in terms of properties?

A
  • similar properties occurred every element
  • law of octaves
  • broke down after calcium
58
Q

What did Mendeleev develop?

A
  • he developed the first modern periodic table and overcame some of the problems of early periodic tables;
59
Q

What did that Mendeleev do that no one had done before in the periodic table?

A
  • he left gaps in is periodic table for elements that he thought had not been discovered;
  • this allowed Mendeleev to predict the undiscovered elements’ properties
  • he changed the order of specific elements
60
Q

How did Mendeleev order the elements in the period table?

A
  • Mendeleev started by arranging all elements in order of increasing atomic mass, but not always strictly
  • i.e. in some places he changed the order based on atomic weights
61
Q

What caused Mendeleev’s periodic table to be accepted?

A
  • because he left gaps in the table which allowed him to predict the undiscovered elements’ properties, when these elements were discovered to have the same properties he predicted
  • his ideas were confirmed and his table was widely accepted
62
Q

What did the discovery of isotopes confirm?

A
  • they explained why the order based on atomic weights was not always correct
  • this is because isotopes of the same element have different masses but the same chemical properties so occupy the same position on the periodic table but can appear too heavy or too light, distorting the order of atomic weights.
  • this confirmed Mendeleev was correct
63
Q

What are the differences between the modern periodic table and Mendeleev’s?

A
  • it is ordered in order of atomic number rather than atomic weight (protons had not been discovered when Mendeleev made his table)
  • the modern periodic table has group 0, the noble gases; these were not fully discovered when Mendeleev published his table
64
Q

Why is relative atomic mass used when referring to elements on the periodic table?

A
  • because elements can have several isotopes, hence why some mass numbers are not whole numbers
65
Q

What do periods (rows) tell you in the periodic table?

A
  • how many electron shells the element has
66
Q

Define an ion

A
  • an atom formed by the loss or gain of electrons, which is now charged
67
Q

What happens when metals react?

A
  • they lose electrons to achieve a full outer shell and become stable;
  • this loss of electrons means that metals react to form positive ions because there are more protons than electrons when an electron is lost
68
Q

Why do metals lose electrons to form ions?

A
  • because they have very few electrons in their outer shell
  • so it is easier to lose electrons to gain a stable electronic structure (full outer shell) than gaining electrons;
  • it requires less energy to lose than gain
69
Q

What forms positive ions and what forms negative ions?

A
  • positive ions -> metals
  • negative ions -> non-metals
70
Q

What are the majority of elements?

A
  • metals
71
Q

Describe the properties of metals

A
  • strong
  • can be bent and hammered into different shapes (malleable)
  • great conductors of heat and electricity
  • high melting and boiling points so usually solid at room temperature (except mercury)
72
Q

Describe the properties of non-metals in comparison to metals

A
  • duller and more brittle than metals
  • lower boiling and melting points so are sometimes gases at room temperature
  • poor conductors of heat and electricity
  • often have a lower density than metals
73
Q

Why are group 0 gases very unreactive?

A
  • because they have a full outer shell of electrons (except helium which has 2)
  • so they do not need to gain or lose any electrons via a reaction, because they have a stable arrangement of electrons
  • therefore they do not easily form molecules
74
Q

Why are noble gases used in lightbulbs?

A
  • because of their stable electronic structure, they’re unreactive,
  • meaning they’re good for use in light bulbs because they’re non-flammable so are less likely to ignite in the heat
75
Q

What happens as you go down group 0?

A
  • relative atomic mass increases. * as this increases, the boiling points increase
76
Q

Group 1 aka
Group 0 aka
Group 7 aka

A
  • 1 - alkali metals
  • 0 - noble gases
  • 7 - halogens
77
Q

Describe how alkali metals react with oxygen

A
  • they react rapidly with oxygen in oxidation reactions to form metal oxides
  • e.g. lithium + oxygen -> lithium oxide
78
Q

Describe how alkali metals elements react with chlorine

A
  • they react rapidly with chlorine to form metal chlorides (a white precipitate)
  • e.g. lithium + chlorine -> lithium chloride
79
Q

Describe how alkali metals react with water and what you would see when adding UI

A
  • they react rapidly and vigorously with water to form metal hydroxides and hydrogen;
  • there is effervescence as hydrogen is produced
  • since metal hydroxides are alkali, if we placed UI in the reaction between alkali metals (and any metal, in fact) and water, it would go from green (pH7) top purple (alkaline pH>7)
80
Q

Describe how reactivity changes when reacting alkali metals with water as you go down the group

A
  • the further down group 1 (the more reactive), the more violent the reaction;
  • hence francium + water is very explosive, but sodium + water is only a little bit of fizzing
81
Q

How should alkali metals be handles and why?

A
  • they should be stored in oil and handled with forceps,
  • because alkali metals are very reactive
82
Q

Why do groups have similar reactions?

A
  • because they all have the same number of electrons in their outer shell
83
Q

Why does the reactivity of alkali metals increase as you go down the group?

A
  • as you move down group 1, the outermost electron is further away from the nucleus, because the atoms get larger and have more shells.
  • this means that there are weaker electrostatic forces of attraction between the negative outer electron and the positive nucleus,
  • so it is easier to lose this electron during a reaction; hence reactivity increases
84
Q

What may weaken electrostatic forces of attraction?

A
  • the shielding effect from the internal energy levels
  • increased distance of electrons from the nucleus
85
Q

What decreases/increases as you go down group 1?

A
  • reactivity increases
  • melting and boiling points decrease
  • increased relative atomic mass
86
Q

What decreases/increases as you go down group 0?

A
  • reactivity stays the same
  • melting and boiling points increase
  • increased relative atomic mass
87
Q

What are halogens?

A
  • non-metals that consist of molecules made of pairs of atoms;
  • they have coloured vapours
88
Q

Why do all halogens form diatomic molecules?

A
  • each atom needs to gain 1 electron to have a stable electronic configuration,
  • so if it bonds to another atom by a covalent bond, both atoms in the molecule become stable
89
Q

Why does boiling and melting point increase as you go down group 7?

A
  • the relative formula mass of the molecules increases
  • due to increased relative formula mass, the molecules get larger meaning there are stronger intermolecular forces which require more energy to overcome
  • because of this, they have increasing melting and boiling points as you go down the group * (hence why top few halogens are gases but bottom few are solid)
90
Q

Why does reactivity decrease as you go down group 7?

A
  • as you go down the group, the outer shell of electrons gets further away from the nucleus.
  • this increased distance, together with the increased shielding effect of inner shells, weakens the electrostatic forces of attraction between the positive nucleus and the outer shell,
  • so it is harder to gain electrons in reactions as you go down the group
91
Q

What do halogens form when they react with metals?

A
  • ionic compounds in which the halide carries a -1 charge
92
Q

What do halogens form when they react with non-metals?

A
  • covalent compounds where there is a shared pair of electrons
93
Q

What happens when you put two different halogens in an aqueous solution?

A
  • the more reactive halogen can displace the less reactive halogen from an aqueous solution of its salt
  • e.g. with sodium bromide + fluorine, fluorine is more reactive than bromine, so will displace it
94
Q

What decreases/increases as you go down group 7?

A
  • reactivity decreases
  • boiling and melting point increase
  • relative atomic mass increases
95
Q

What are the transition metals?

A
  • metals with similar properties which are different from those of the elements in group 1
96
Q

Compare melting points of transition metals vs alkali metals

A
  • group 1 metals have relatively low melting points,
  • whilst transition metals have relatively high boiling points (apart from mercury)
97
Q

Compare densities of transition metals vs alkali metals

A
  • group 1 metals have a relatively low density (Li, Na and K are less dense than water)
  • whereas transition metals have a high density
98
Q

Compare strength of transition metals vs alkali metals

A
  • group 1 metals are soft and can be cut with a knife,
  • whereas transition metals are hard and strong e.g. iron
99
Q

Compare reactivity with oxygen, halogens and water of transition metals vs alkali metals

A
  • group 1 metals all react very rapidly with oxygen halogens and water,
  • whereas transition metals are much less reactive than group 1 metals
100
Q

Compare ions of transition metals vs alkali metals

A
  • transition metals can form ions with different charges
  • e.g., iron can form Fe2+ or Fe3+, * but group 1 metals can only form 1+ ions
101
Q

Describe the typical properties of a transition metal

A
  • they have ions with different charges
  • they can form coloured compounds
  • they are useful as catalysts
102
Q

Describe lithium’s reaction with oxygen, water, and chlorine

A
  • oxygen - burns with a red flame and produces a white solid
  • water - fizzes steadily, gradually disappears
  • chlorine - white powder is produced and settles on the sides of the container
103
Q

Describe sodium’s reaction with oxygen, water, and chlorine

A
  • oxygen - orange flame produces white solid
  • water - fizzes rapidly, melts into ball and disappears quickly
  • chlorine - burns with yellow flame, clouds of white powder produced and settles on side of container
104
Q

Describe potassium’s reaction with oxygen, water, and chorine

A
  • oxygen - large pieces produce lilac flame, smaller makes solid immediately
  • water - ignites with sparks and lilac flame, disappears quickly
  • chlorine - more vigorous than sodium

ZZ Chem 1 (5 decks)

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