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Chemistry > Elements of the periodic table > Flashcards

Elements of the periodic table Flashcards

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1
Q

What is nanotechnology? what are examples?

A
  • the science and technology of the manipulation of the structure of matter at a molecular level.
  • Nanoechnology is concerned with very small things measured in nanometers.
  • Sunscreen, computers and mobile phones utilise this and it could help with disease treatment and computer efficiency.
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2
Q

what is the nanometre?

A

1 billionth of a metre or 10 to the power of -9. It is the size of six carbon atoms.

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3
Q

what does the nanoscale describe?

A

things about 1-100nm wide (eg. DNA 2nm wide)

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4
Q

how do you convert to nanometers?

A

convert to meters and then to nanometers by multiplying by either 10^9 or ^-9

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5
Q

how can a material be classified as a nonmaterial?Examples? How are they different to materials in bulk?

A
  • they must have one dimension less than 100nm eg. nano layers, nanowires or nanotubes
  • they have different properties
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6
Q

what is a nanoparticle?

A

a type of nonmaterial that is usually spherical and has a diameter less than 100nm

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7
Q

how can nanoparticles be used?

A

to catalyse, the surface can adsorb (attach) to different molecules and transport molecules around the body, can be used in medicine, cancer treatment, computer storage etc.

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8
Q

what are risks of nanoparticles?

A

they can travel through skin into the bloodstream and into cells. they can cause unwanted chemical reactions in the body and no one knows what they do and the long term effects

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9
Q

what is atomic theory?

A

the best explanation but it can change and develop

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10
Q

who presented the first atomic theory? when? what was it?

A

in 1802, John Dalton.

  • He proposed that all matter was made up of small spherical particles that are invincible and indestructible.
  • elements are made of one type of atom and compounds, many to fixed ratios
  • elements could not be broken down
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11
Q

how was dalton wrong?

A

that elements could not be broken down. WE now know they have subatomic particles

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12
Q

what is a molecule?

A

when non-metals combine together. many gases are diatomic molecules (O2)

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13
Q

what are monatomic atoms?

A

noble gases exist as individual atoms because they are chemically inert

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14
Q

when is a molecule known as as a compound?

A

if it is made up of two or more different elements

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15
Q

how many elements occur naturally?

A

around 92

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16
Q

what 5 ways can elements get symbols?

A

first letter, first letter and another letter, first two letters, Latin name, from other languages

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17
Q

nucleus

A

made up of protons and neutrons (nucleons)

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18
Q

protons

A

positive

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19
Q

neutrons

A

neural

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20
Q

electrons

A

negative

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21
Q

how much lighter is an electron than a proton or neutron?

A

1800 times

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22
Q

how much of the atoms mass does the nucleus take up?

A

99.97%

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23
Q

what did Ernest Rutherford propose and when? (4)

A
  • most mass and all + charge is located in a central region
  • most volume is empty space occupies by electrons
  • electrons circle the nucleus
  • the force between the nucleus and electrons is electrostatic
  • 1911
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24
Q

What was Rutherford’s experiment to prove? what did he do? What did he conclude?

A
  • wanted to prove Thomson’s plumb pudding model correct
  • fired alpha particles (+) from a radium at thin gold foil
  • the paths of particles were determined by bursts of light produced on the screen behind the foil
  • most particles went through
  • 1 in 8000 deflected and one bounced back
  • conc: most of an atom is empty space and the alpha particles bounce back when repelled by positive matter
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25
Q

What did James Chadwick discover? What did he propose? when?

A
  • identified the neutron

- proposed that the nucleus also contained neutrons which explained isotopes

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26
Q

what is an isotope? properties? example?

A
  • same number of protons but differing number of neutrons
  • they have identical chemical properties but different physical properties (mass and density)
  • every atom has isotopes
  • C-14 breaks down into another atom (radioactive)
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27
Q

what is an ion? what are the different types?

A
  • result of atoms gaining or losing electrons
  • cation +
  • anion -
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28
Q

what colour dow potassium produce?

A

lilac

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29
Q

what colour does lithium produce?

A

red

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30
Q

what colour does barium produce?

A

yellow-green

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31
Q

what colour does calcium produce?

A

orange-red

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32
Q

what colour does copper produce?

A

blue-green

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33
Q

what colour does strontium produce?

A

scarlet

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34
Q

what colour does sodium produce?

A

orange

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35
Q

why do some metal ions produce a colour if heated?

A
  • when heated electrons can gain energy and jump to higher energy levels
  • this makes them energetically unstable so they fall back, releasing light energy
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36
Q

do all atoms produce electromagnetic energy if heated? how can it be seen?

A

yes. if the light is passed through a prism onto a screen, an emission spectrum is obtained

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37
Q

what is a spectrum?

A

a series of lines on a black background

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38
Q

What are emission spectra?

A

they are related to the structure within an atom and each one has a unique spectra

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39
Q

what did Bohr propose and when?

A
  • that electrons circled the nucleus in fixed shells
  • shells closer to nucleus have lower energy and those furthest have high energy
  • electrons move in orbitals
  • heating an element can cause can cause the electron to gain energy and jump to another level
  • it then returns to ground state, releasing energy
  • they can do this in a number of ways and this produces different colours
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40
Q

excited state

A

when electrons have jumped to a higher energy state

41
Q

ground state

A

when are electrons are In their lowest energy state

42
Q

what do electrons in the same shell share in common?

A

the same amount of energy

43
Q

what are shells more like?

A

clouds

44
Q

do electrons closer or further from the nucleus have more attraction?

A

closer

45
Q

when is it easiest to remove an electron? what is this amount of energy called?

A

the further away the electron is. ionisation energy

46
Q

what are the rules for shells?

A
  • each shell holds a max number of electrons
  • lower energy shells fill before higher ones
  • electrons shells fill in an order
47
Q

what is the equation for the max number of electrons in a shell?

A

2n^2

48
Q

what electrons are involved in reactions?

A

valance electrons

49
Q

only eight electrons can be in the last shell. what is tis rule?

A

octet rule

50
Q

what is a major rule in the filling of shells?

A

put 8 in the third shell, then 2 in the 4th shell and then fill up the third shell

51
Q

what did Schrodinger discover in 1926?

A
  • he developed a quantum mechanical model where electrons occupied regions of space called orbitals where the probability of finding an electron was high.
  • electrons move in major levels called shells and within these are sub shells with regions given different letters
52
Q

what does the Pauli exclusion principle say?

A

that an orbital can have 0, 1, or 2 electrons

53
Q

how many can the s sub shell hold?

A

1 orbital so 2 electrons

54
Q

how many can the p sub shell hold?

A

3 orbital so 6

55
Q

how many can the d subshell hold?

A

5 orbitals so 10

56
Q

how many can the f sub shell hold?

A

7 orbitals so 14

57
Q

how do you write notation in the exited state?

A

an electron may be out of place and in another sub shell

58
Q

how is notation written with negative ions?

A

number of electrons is determined and then they are filled normally

59
Q

how is notation written with positive ions?

A

the notation is written for a neutral atom and then electrons are removed from the sub shell furthest from the nucleus

60
Q

what atoms don’t follow the rules?

A
  • chromium: a half-filled 3d sub shell is preferred for stability
  • copper: filled 3d shell is more stable
61
Q

who started purifying, preparing, identifying, assigning letters and calculating masses of elements and compounds?

A
  • Berzelius
62
Q

what did Mendeleev do? when? how did he organise this?

predictions?

A
  • created the periodic table in 1869
  • ordered in mass order
  • organised chemical data by properties and concluded that chemical properties varied periodically with increasing mass
  • elements with similar properties were grouped in groups
  • left gaps for things undiscovered such as geranium which he predicted the grey colour and high melting point correctly
63
Q

how do elements vary?

A

by their electronic configuration

64
Q

what is periodicity?

A

the repeated pattern of similar chemical properties as elements are arranged in order of increasing atomic number

65
Q

what do elements in the same period share?

A

the same number of shells

66
Q

what do elements in the same group share?

A

same number of electrons in outer shell

67
Q

what is group 1?

A

alkali metals

68
Q

group 2?

A

alkaline earth metals

69
Q

what does it mean if the last sub shell to be filled or partially filled is the is a d-sub Shell? then the element is a transition metal and it is usually not given a group number

A

then the element is a transition metal and it is usually not given a group number

70
Q

what is the s-block?

A
  • groups 1 and 2
  • s orbital is highest energy sub shell
  • electronic configuration ranges from s1 to s2
71
Q

what is the p-block?

A
  • groups 13-18
  • p orbital is highest energy sub shell
  • outer shell configurations range from s1p1 to s2p6
72
Q

what is the d-block?

A

contains transition metals. Elements have orbitals in which d-subshells are progressively being filled. D-subshells are only filled after the s-subshell of the next shell is filled. The outer shell configurations range from, d1s2 to d10s2

73
Q

what is the f-block?

A

lanthanides and actinides. In the lanthanides, the orbitals in 4f are progressively being filled and in the actinides it is the orbitals in 5f

74
Q

what can core charge explain?

A

Atomic radii, electronegativity and ionisation energy

75
Q

what is core charge?

A

the attraction felt by the valance electrons to the positive nuclear charge.

76
Q

explain the other force (other than attraction to nucleus) acting on the valance electrons?

A

Valance electrons are held to the nucleus but are also shielded by other electrons which they are repelled by.

77
Q

how can core charge be calculated?

A

number of valance electrons

78
Q

what are the trends in core charge down a group?

A

Core charge stays constant down a group, but valance electrons are held less strongly as they are further from the nucleus (more shells in the atom)

79
Q

what are the trends in core charge across a period?

A

Core charge increases from left to right across a period. The valance electrons are more attracted to the nucleus as the core charge increases.

80
Q

what is electronegativity?

A

the ability of an atom to attract electrons towards itself.

81
Q

how does electronegativity relate to core charge?

A

The greater the core charge, the greater the electronegativity.

82
Q

what has the highest electronegativity?

A

fluorine (4.0)

83
Q

what is the trend in electronegativity down a group?

A

Down a group core charge stays constant and the number of shells increases. Therefore valence electrons are less strongly attracted to the nucleus as they are further from the nucleus. As a result, electronegativity decreases.

84
Q

what is the trend in electronegativity across a period?

A

Left to right across a period, electronegativity increases. The number of shells remains constant but the core charge increases across a period. Thus, the valance electrons become more strongly attracted to the nucleus. As a result, electronegativity increases.

85
Q

what is atomic radius?

A

distance from the nucleus to the valance electrons.

86
Q

what is the trend in atomic radius down a group?

A

As we go down a group, the core charge stays constant and the number of shells increases down a group. Therefore atomic radii increase.

87
Q

what is the trend in atomic radius across a period?

A

From left to right across a period, the number of occupied shells remains constant but the core charge increases. Thus, the valance electrons become more strongly attracted to the nucleus and as a result, the atomic radius decreases.

88
Q

what is ionisation?

A

The process of removing electrons to form an ion is called ionisation.

89
Q

which electrons are removed first?

A

valance electrons because they are less tightly held

90
Q

what is the energy required to remove the first electron from an atom? what does this energy depend on?

A

The energy required to remove the first electron from an atom is called the first ionisation energy.
The size of this energy depends on how tightly held the electron is.

91
Q

what is the trend in ionisation energy down a group?

A

Down a group the first ionisation energy decreases because core charge stays the same and the number of shells increases down a group. This means that electrons are less tightly held because they are further away from the nucleus. As a result less energy is needed to remove the electron so ionisation energy decreases.

92
Q

what is the trend in ionisation energy across a period?

A

From left to right across a period, this increases because the number of shells remains constant but the core charge increases across the period. trus, the valance electrons are more strongly attached and more energy is required to remove them.

93
Q

percentages of metals and non-metals?

A
  • 75% are metals and are on the left of the periodic table
  • 20% are non-metals and they are on the right
  • Some have both qualities, these are metalloids
94
Q

what is the trend in metal reactivity down a group?

A

Down a group, the core charge is constant and the number of shells is greater so the atomic radii is greater. So the electrons are less tightly held. As a result they are more easily lost and the element is more reactive.

95
Q

what is the trend in metal reactivity across a period?

A

Across a period, the core charge is greater and the number of shells is constant so the atomic radii is smaller. The electrons are held more tightly held and as a result are harder to remove, so the metal is less reactive.

96
Q

what does metal reactivity depend on?

A

how easily electrons can be removed

97
Q

what is the trend in reactivity of non-metals down a group?

A

Down a group, the core charge is constant and the number of shells increases so the atomic radius is larger. This means that the valance shell is further from the nucleus and as a result, electrons are less attracted and the non-metal is less reactive.

98
Q

what is the trend in reactivity of non-metals across a period?

A

Across a period, the core charge is larger and the number of occupied shells is the same so the radius is smaller. Hence, the outermost shell is closer to the nucleus and as a result, electrons are more easily attracted so the non-metal is more reactive.

99
Q

do noble gases react? why not?

A

no, Noble gases are inert. This is because they have a stable electronic structure, with 8 electrons in their outer shell except for helium which only needs 2.

Chemistry (28 decks)

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