Chapter 7 Flashcards
Mendeleev
Arranged elements in order of atomic mass
Aligned elements in groups with similar properties
Swapped elements and left gaps
Predicted the properties of missing elements
What is a group
The elements are arranged in vertical columns
Same number of outer shell electrons
What is a period
The elements are arranged in horizontal rows
Number of the highest energy electron shell in the elements atom
Periodicity
A repeating trend in properties of the elements across each period
Trends across a period
Each period starts with an electron in a new highest energy shell
S subshell fills with 2 electrons
P fills with 6
Group 4
3D subshell fills with 5
Trends down a group
Elements in each group have atoms with the same number of electrons in each sub shell
Same chemical properties
Blocks
Divides elements corresponding to their highest energy sub shell
Ionisation energy
Measures how easily an atom loses electrons to form positive ions
First ionisation energy
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
Factors that affect ionisation energy
Attraction between the nucleus and the outer electrons of an atom
Atomic radius
Electron shielding
Nuclear charge
Atomic radius
The greater the distance
between the nucleus and outer electrons
the less the nuclear attraction
the easier it is to lose an electron
Nuclear charge
The more protons there are
Te greater the attraction between the nucleus and outer electrons
Harder to lose electrons
Electron shielding
Inner shell electrons repel outer shell electrons = shielding effect
Reduces the attraction between the nucleus and outer electrons
Why is the second ionisation energy is greater than the first
Removed from positive ion, more protons than electron
Greater pull on electrons from nucleus
Less shielding and repulsion from other electrons so held more closely
Large jump in ionisation energy across shells because - less shielding
What predictions can you make from successive ionisation energies
The number of electrons in the outer shell
The group of the element
The identity of the element
Trends in first ionisation energies of elements across a period
General increase in first ionisation energy
Sharp decrease between the end of one period and the start of another
Trend in first ionisation energy down a group
FIE decreases down a group
Atomic radius increases
Shielding increased
Nuclear attraction on outer electrons decreases
Trend in first ionisation energy across a period
General increase of FIE
Nuclear charge increases
Same shell - same shielding
Nuclear radius decreases
Nuclear attraction increases
Subshell trends in first ionisation energy
First drop - 2,3
The start of filling the 2p subshell
2p subshell in boron has a higher energy than the 2s subshell in beryllium
Therefore 2p electron is easier to remove than one of the 2s electrons in beryllium
Why does FIE rise and fall across a period
Due to the existence of subshells, their energies, and how they fill with electrons
Subshell trends in first ionisation energy
Second drop - 5,6
Phosphorous has 3 electrons in the 3p subshell (1 in each)
Sulfur has 4 electrons in the 3p subshell (2 paired in one orbital, 1 in each of the other 2)
The paired electrons in sulfur repel each other
Making it easier to remove one of these electrons than an unpaired one
Metalloids
Elements near the metal/non metal divide
Metallic bonding and structure
Strong electrostatic attraction between positive cations and negative delocalised electrons
Billions of metal atoms held by metallic bonding to form giant metallic lattices
Properties of metals
Strong metallic bonds
High electrical conductivity
High mp bp
Properties of giant covalent lattice structures
High mp,bp
Insoluble - covalent bonds can’t be broken by interactions with solvents
Electrical conductivity - none except graphene and graphite
Graphene and graphite
Carbon based on planar hexagonal layers - 120
Each carbon only forms 3 bonds
Spare electrons are delocalised
Periodic trend in melting points
Melting point increases from 1-14
As metallic bonds get stronger
Because metallic ions have an increasing positive charge, more delocalised electrons, smaller atomic radius
Giant covalent lattices
Structure made of strong covalent bonds that link atoms together
A lot of energy needed to break these bonds
Molecular substances
Mp depends on strength of London forces
London forces are weak
Larger atoms with more electrons have stronger London forces
Argon is low because monoatomic
Allotrope
Different forms of the same element in the same
Structure of graphite
Carbons arranged in sheets of hexagons
Each carbon covalently bonded to 3 other carbons
4th outer electron is delocalised
Layers of sheets held together by London forces
Similarity and differences between first ionisation across period 2 and 3
Similar
Same trend
Because similar subshell structure across period
Difference
Period 3 would have lower ionisation values
Because electrons are further away from nucleus and experience more shielding
