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Chemistry > M3 Periodicity > Flashcards

M3 Periodicity Flashcards

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1
Q

Features of Mendeleev’s periodic table

A
  • Organised by atomic mass
  • Blank spaces left for elements he predicted hadn’t been discovered
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2
Q

Features of the modern periodic table

A
  • Organised by atomic number
  • Blank spaces filled in with more recent discoveries
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3
Q

Features of both periodic tables

A
  • Organised elements in periods and groups
  • Elements organised in groups with similar chemical properties/trends in physical properties
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4
Q

How does atomic number vary on the periodic table?

A

Reading from left to right, elements are arranged in order of increasing atomic number. Each successive element has atoms with one extra proton.

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5
Q

Describe groups in the periodic table

A

Elements are arranged in vertical columns, each element has atoms with the same number of outer-shell electrons and similar properties.

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6
Q

Describe periods in the periodic table

A

Elements are arranged in horizontal rows called periods. The number of the period gives the number of the highest energy electron shell in an elements atoms.

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7
Q

What is periodicity?

A
  • A repeating trend in properties of the elements.
  • Trends across periods include electron configuration, ionisation energy, structure and melting points.
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8
Q

Describe periodic trends in electron configuration

A
  • Across period 2, the 2s sub-shells fills with two electrons, followed by the 2p sub-shell with 6 electrons.
  • Across period 3, the same pattern of filling is repeated for the 3s and 3p sub-shells.
    Period 1 and 2 = s-block
    Transition metals = d-block
    Periods 3 - 8 = p-block
    Two rows below = f-block
  • Elements in each group have atoms with the same number of electrons in each sub-shell
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9
Q

What is the name of group 1

A

Alkali metals

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10
Q

What is the name of group 2

A

Alkaline each metals

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11
Q

What is the name of group 3 - 12 (new)

A

Transition metals

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12
Q

What is the name of group 7/17

A

Halogens

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13
Q

What is the name of group 0/18

A

Noble gases

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14
Q

What is ionisation energy?

A

Ionisation energies measure how easily electrons can be lost from an atom to form positive ions.

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15
Q

What is the first ionisation energy?

A

The energy required to remove one electron from each atom in one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions.

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16
Q

What are the factors affecting ionisation energy?

A
  • Electrons are held in their shells by attraction from the nucleus. The first electron lost will be in the highest energy level and will experience the least attraction from the nucleus.
  • The factors affecting attraction between the nucleus and the outer electrons is atomic radius, nuclear charge and electron shielding.
17
Q

How does atomic radius affect ionisation energy?

A

The greater the distance between the nucleolus and the outer electrons the weaker the nuclear attraction. I’m he force of attraction decreases rapidly as the distance increases, hence atomic radius has a large effect on ionisation energy.

18
Q

How does nuclear charge effect ionisation energy?

A

The greater the number of protons in the nucleus of an atom, the greater the attraction between the nucleus and the outer electrons.

19
Q

How does electron shielding effect the ionisation energy?

A

Electrons are negatively charged and so inner-shell electrons repel outer-shell electrons. This repulsion, called the shielding effect reduces the attraction between the nucleus and the outer electrons.

20
Q

How many ionisation energies can an element have?

A

An element has as many ionisation energies as there are electrons.
He (g) = He+(g) + e- (first ionisation energy)
He+(g) = He2+(g) + e- (second ionisation energy)
- The second ionisation energy of helium is greater than the first ionisation energy. There are two protons attracting two electrons in the 1s sub-shell. After the first electron is lost, the single electron is pulled closer to the helium nucleus. The nuclear attraction on the remaining electron increases and more ionisation energy will be needed to remove the second electron.

21
Q

What can cause large differences in ionisation energy?

A
  • A large increase between ionisation energies suggests that the next electron must be removed from a different shell, making it closer/further away and with less/more shielding.
  • Or when an electrons pair in orbitals, so the ionisation energy may be lowered for the first element with electrons pairing in orbitals due to electron repulsion.
22
Q

What are the patterns in first ionisation energies in the first 20 elements?

A
  • There is a general increase in first ionisation energy across each period
  • There is a sharp decrease in first ionisation energy between the end of one period and the start of the next period
23
Q

What are the trends in first ionisation energy down a group?

A
  • First ionisation energies decrease down a group.
  • This is because although nuclear charge increases, it’s effect is outweighed by the increased radius and (but to a lesser extent) the increased shielding.
  • Down a group the atomic radius and number of inner shells increase so shielding increases. Nuclear attraction on outer electrons decreases and first ionisation energy decreases.
24
Q

What are the trends in first ionisation energy across a period?

A
  • Across a period nuclear charge increases, as it is the same shell the shielding stays similar, nuclear attraction increases, atomic radius decreases and first ionisation energy increases.
25
Q

What are sub-shell trends in first ionisation energy?

A
  • First ionisation energy shows a general increase across period 2 and 3, however it does fall in two places, the drops occur at the same position in each period.
  • In period 2 the drop from Be to B marks the filling of the 2p sub-shell. The 2p sub-shell in boron has a higher energy than the 2s sub-shell in beryllium, therefore in boron the 2p electron is easier to remove than one of the 2s electrons in beryllium. The first ionisation energy of boron is less than the first ionisation energy of beryllium.
  • The fall in first ionisation energy from nitrogen to oxygen marks the start of electron pairing in the p-orbitals of the 2p sub-shell. In nitrogen and oxygen the highest energy electrons are in a 2p sub-shell. In oxygen the paired electrons in one of the 2p orbitals repel one another, making it easier to remove an electron from an oxygen atom than a nitrogen atom, therefore the first ionisation energy of oxygen is less than the first ionisation energy of nitrogen.
    Three 2p electrons: one spin in each 2p orbital, spins are at right angles due to equal repulsion between electrons
    Four 2p electrons: two electrons in one 2p orbital, 2p electrons start to pair and the paired electrons repel
26
Q

Define and describe metallic bonding

A

Metallic bonding is the strong electrostatic attraction between cations and delocalised electrons.

  • In a solid metal structure, each atom has donated its negative outer-shell electrons to a shared pool of electrons, which are delocalised throughout the whole structure.
  • The positive ions (cations) left behind consist of the nucleus and the inner electron shells of the metal atoms.
  • The cations are fixed in position, maintaining the structure and shape of the metal.
  • The delocalised electrons are mobile and able to move throughout the structure.

In a metal structure, billions of metal atoms are held together metallic bonding in a giant metallic lattice.

27
Q

What are the properties of metals?

A
  • Strong metallic bonds (attraction between positive ions and delocalised electrons)
  • High electrical conductivity
  • High melting and boiling points
28
Q

Describe the electrical conductivity of metals

A
  • Metals conduct in solid and liquid states.
  • When voltage is applied across a metal the delocalised electrons move through the structure, carrying the charge.
29
Q

Describe the melting and boiling points of metals

A
  • Most metals have high melting and boiling points.
  • The melting point depends on the strength of the metallic bonds holding together the atoms in the giant metallic lattice:
    • For most metals, high temperatures are necessary to provide the large amount of energy needed to overcome the strong electrostatic attraction between cations and electrons.
    • This strong attraction results in most metals having high melting and boiling points.
30
Q

Describe the solubility of metals

A
  • Metals do not dissolve.
  • Any interactions between polar solvents and the charges in a metallic lattice lead to a reaction, rather than dissolving.
31
Q

Describe giant covalent structures

A
  • Many non-metallic elements exist as simple covalently bonded molecules. In a solid state these molecules form a simple molecular lattice held together by weak intermolecular forces.
  • These structures therefore have low melting and boiling points.
  • The non-metals boron, carbon and silicon have different lattice structures, instead of small molecules and intermolecular forces many billions of atoms are held together by a network of strong covalent bonds to form a giant covalent lattice.
  • Carbon and silicon’s atoms have four electrons in the outer shells, they use these to form covalent bonds to other carbon or silicon atoms, resulting in a tetrahedral structure (the bond angles are all 109.5° due to electron-pair repulsion)
32
Q

Describe the melting and boiling points of giant covalent structures

A
  • Giant covalent lattices have high melting and boiling points, due to strong covalent bonds.
  • High temperatures are necessary to provide the large amount of energy needed to break the strong covalent bonds.
33
Q

Describe the solubility of giant covalent lattices

A
  • Giant covalent lattices are insoluble in almost all solvents.
  • The covalent bonds holding together the atoms in the lattice are far too strong to be broken by interactions with solvents.
34
Q

Describe the electrical conductivity of giant covalent lattices

A
  • Giant covalent lattices are non-conductors of electricity. The only exceptions are graphene and graphite, which are forms of carbon.
  • In carbon (diamond) and silicon all four outer-shell electrons are involved in covalent bonding, so none are available for conducting electricity.
  • Carbon is special in forming several structures in which one of the electrons is available for conductivity (graphene and graphite)
35
Q

Why can graphene and graphite conduct electricity?

A
  • Only three electrons of the four outer-shell electrons are used in covalent bonding. The remaining electron is released into a pool of delocalised electrons shared by all atoms in the structure. Structures of carbon containing planar hexagonal layers are therefore good electrical conductors.
  • Graphene and graphite are both giant covalent structure of carbon based on planar hexagonal layers with bond angles of 120° by electron-pair repulsion.
36
Q

Describe the structure of graphene

A
  • Graphene is a single layer of graphite, composed of hexagonally arranged carbon atoms linked by strong covalent bonds.
  • Graphene has the same electrical conductivity as copper, and is the thinnest and strongest material ever made.
37
Q

Describe the structure of graphite

A
  • Graphite is composed of parallel layers of hexagonally arranged carbon atoms.
  • The layers are bonded by weak London forces.
  • The bonding in the hexagonal layers only uses three of carbons four outer electrons, the spare electron is delocalised between the layers so electricity can be conducted as in metals.
38
Q

Describe the periodic trend in melting points across period 2 and 3

A
  • The melting point increases from Group 1 to Group 14 (4)
  • There is a sharp decrease in melting and point between Group 14 and 15 (4 and 5)
  • The melting points are comparatively low from Group 15 (5) to Group 18 (0)

• The sharp increase in melting point marks a change from giant to simple molecular structures.
• On melting, giant structures have strong forces to overcome so have high melting points. Simple molecular structures have weak forces to overcome, so have much lower melting points.
• This trend in melting points across period 2 is repeated across period 3 and continues across the s- and p-blocks from period 4 downwards.

Chemistry (15 decks)

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