Unit 1: KA1 Periodicity Flashcards
The periodic table arranges all chemical elements in extraordinary ways. Different types of chemical bonding, patterns and trends can be observed in their arrangement. This deck of flashcards covers everything in Key Area 1 (Periodicity) of Higher Chemistry: groups, periods, types of bonding and structure, covalent radii, ionisation energies, electronegativities and every past paper question (except the repetitive ones) on Periodicity. Hopefully, it'll help you just as much as it helped me. :))
How are elements arranged in the periodic table?
In order of increasing atomic number.
What does the periodic table allow chemists to do?
Make accurate predictions of physical properties and chemical behaviour for any element, based on its position.
What are features of the periodic table (arrangement of elements-wise)?
Groups
Periods
GROUPS in the periodic table?
Vertical columns within the table.
What do “Groups” contain?
Contain elements with similar chemical properties, resulting from a common number of electrons in the outer shell.
PERIODS in the periodic table?
Rows of elements arranged in increasing atomic number.
What do “Periods” demonstrate?
Demonstrates an increasing number of outer electrons,
Moves from metallic to non-metallic characteristics.
How are the first 20 elements in the periodic table categorised?
According to bonding and structure.
Name the categories that the first 20 elements of the periodic table are sorted into. [4 answers]
Metallic,
Covalent molecular,
Covalent network,
Monatomic
Metallic bonding occurs between atoms of —– elements.
The outer electrons are ———– (free to move).
This produces an ————- force of attraction between the positive —– —- and the negative ———– electrons.
This ———– ‘sea of ———’ allow metals to conduct electricity.
metal,
delocalised,
electrostatic,
metal ions,
delocalised,
delocalised,
electrons
Give examples of METALLIC bonding/structure.
Li,
Be,
Na,
Mg,
Al,
K,
Ca
Describe the COVALENT BOND.
When two POSITIVE NUCLEI (of two non-metal atoms) are held together by their common ELECTROSTATIC ATTRACTION FOR a shared pair of ELECTRONS.
COVALENT MOLECULES are small groups of atoms held together by STRONG ——– bonds INSIDE the molecule and WEAK intermolecular* forces BETWEEN the molecules.
COVALENT.
*Refer to KA2 for the
weak London Dispersion forces between molecules.
Give examples of COVALENT MOLECULAR bonding/structure.
Most of the discrete covalent molecules are diatomic elements:
H₂,
N₂,
O₂,
F₂,
Cl₂,
There are also larger covalent molecular elements:
P₄,
S₈,
Fullerenes (eg. C₆₀)
Why are most of the diatomic molecules GASES at room temperature?
The diatomic molecules have very small intermolecular (London dispersion) forces. There’s enough energy at room temperature to overcome these weak forces of attraction, hence the molecules separate from each other (resulting in gases).
Why are Bromine and Iodine (diatomic covalent molecules) solid and liquid at room temperature?
They’re bigger molecules –> have more electrons than smaller gas molecules.
Therefore the London dispersion (intermolecular) forces are much stronger between their molecules –> takes more energy (than available at room temperature) to separate them.
Explain why Sulfur has a higher melting point than Phosphorus.
Sulfur consists of S₈ molecules and phosphorus P₄. The sulfur molecules have more electrons than the phosphorus molecules and so have stronger intermolecular (London dispersion forces). This in turn means sulfur has a higher melting point than phosphorus.
What are MOLECULAR FORMS of Carbon?
Fullerenes (eg. C₆₀)
What are FULLERENES comprised of?
What can fullerenes also exist as?
A series of carbon atoms arranged in hexagons and pentagons.
These are joined to make a ball shape.
They can also exist as tube shapes called nanotubes.
Fullerenes are very large molecules and so lots of energy is required to separate them. True or False?
True
Fullerenes sublime (change from solid to gas) at around 600*C.
Covalent networks are large, rigid 3D arrangements of atoms held together by weak intermolecular forces. True or False?
False, covalent networks are held together by STRONG COVALENT BONDS.
Does it take lots of energy to separate the atoms in a covalent network?
Yes. Lots of energy is required to break the strong covalent bonds.
This explains why elements with covalent network structures have HIGH melting and boiling points.
Give examples of COVALENT NETWORK bonding/structure.
B,
C (diamond & graphite),
Si
Why doesn’t diamond conduct electricity?
In diamond, all the outer electrons of each carbon are used to make single covalent bonds with neighbouring carbon atoms (resulting in a giant covalent network).
There aren’t any unbonded (delocalised) electrons in the diamond structure and so diamond doesn’t conduct electricity.
Why does graphite conduct electricity?
In graphite, each carbon atom is bonded to only 3 other carbon atoms. The 4th electron in each carbon atom is delocalised, so graphite can conduct electricity.
1) Monatomic elements are unstable atoms. True or False?
2) They exist as single, unattached particles. True or False?
3) Do monatomic elements have high or low melting and boiling points?
1) False, they are stable atoms with full electron shells –> so they’re unreactive.
2) True
3) Monatomic elements have LOW MELTING and BOILING POINTS as they’re easily separated by overcoming the weak forces of attraction (London dispersion forces) between the atoms.
Give examples of MON-ATOMIC bonding/structure.
Noble gases - He, Ne, Ar, Kr, Xe, Rn
COVALENT RADIUS?
A measure of the size of an atom.
*HALF THE DISTANCE BETWEEN the NUCLEI of TWO ATOMS joined by a covalent bond.
The trends in the COVALENT RADIUS, across periods and down groups, can be explained in terms of what?
Number of occupied shells, and the nuclear charge.
What does nuclear charge mean?
In this case, it means the pull of electrons to the positive nucleus of an atom.
FIRST IONISATION ENERGY?
How would you write an equation for the first ionisation energy of an element?
Energy required to remove 1 mole of electrons from 1 mole of gaseous atoms.
E(g) –> E⁺(g) + e⁻ (where E stands for ANY ELEMENT including diatomic)
*Always write state symbols!
Write an equation for the first ionisation energy of Helium.
He(g) –> He⁺(g) + e⁻
SECOND (or subsequent) IONISATION ENERGIES?
How would you write an equation for the SECOND ionisation energy of an element?
Refers to the energies required to remove further moles of electrons.
E⁺(g) –> E²⁺(g) + e⁻ (where E stands for ANY ELEMENT including diatomic)
*Always write state symbols!
Write an equation for the second ionisation energy of carbon.
C⁺(g) –> C²⁺(g) + e⁻
Write an equation for the third ionisation energy for Krypton.
Kr²⁺(g) –> Kr³⁺(g) + e⁻
The trends in IONISATION ENERGIES, across periods and down groups, can be explained in terms of what?
Atomic size,
Nuclear charge, and…
Screening effect due to inner shell electrons.
Atoms of different elements have the same attraction for bonding electrons. True or False?
No, different element atoms have different attractions for bonding electrons.
ELECTRONEGATIVITY?
Measure of the attraction an atom involved in a bond has for the electrons of the bond.
The trends in ELECTRONEGATIVITY, across periods and down groups, can be rationalised in terms of what?
Covalent radius,
Nuclear charge, and…
Screening effect due to inner shell electrons.
In the periodic table, elements are arranged in order of decreasing atomic number. True or False?
False, arranged in INCREASING atomic number.
Elements in a (vertical group) do not share similar chemical properties.
False, elements in a group DO SHARE similar chemical properties.
ALKALI METALS?
In Group 1,
Are reactive elements,
Reactivity INCREASES down the group.
HALOGENS?
In Group 7,
Are reactive elements,
Reactivity DECREASES down the group.
NOBLE GASES?
In Group 0.
Are unreactive elements.
Exist as individual atoms.
TRANSITION METALS
In middle section of the Periodic Table (between Groups 2 & 3).
Periods are arranged in order of…
Increasing atomic number.
What is the atomic number?
Number of Protons
Increasing number of protons leads to…
An increasing number of electrons in the outer shell.
As the number of electrons in the outer shell increases, what happens (in terms of the elements)?
The elements move from METALLIC characteristics to NON-METALLIC characteristics.
Atoms increase in size UP or DOWN a group?
DOWN
There is an additional shell of electrons down a —–.
Hence…
Group,
Hence atoms INCREASE in size.
Covalent radius INCREASES or DECREASES down a group?
INCREASES (as atomic size increases)
Atoms INCREASE or DECREASE in size across a period?
DECREASE
Across a period, the same shell of ——— is being filled up.
There is greater positive —— in the ——-.
This pulls the ——– —— TOWARDS the nucleus.
The outer —– is pulled in.
Hence…
electrons,
charge,
nucleus,
electron shells,
shell,
Hence atoms DECREASE in size across period.
Covalent radius INCREASES or DECREASES across a period?
DECREASES
What is IONISATION?
Process of removing electrons from gaseous atoms.
Each ionisation energy is the removal of 1 MOLE of electrons.
Any equations which involve the removal of MORE THAN 1MOL of electrons, requires ionisation energies to be ADDED or MULTIPLIED to calculate total energy required?
ADDED together.
*Basically, more than 1mol of electrons? Add together ionisation energies to get total energy required to remove all the moles.
What happens to the ionisation energy DOWN A GROUP?
Ionisation energy decreases - easier to remove electrons.
Atoms get —— down a group AS…
The outer ——— are ——- from the nucleus.
The outer ——— are also shielded from the full effect of the ——- by the —– ——– ——.
The outer ——— are EASIER or HARDER to remove.
Hence…
bigger,
as there is an additional shell of electrons each time.
electrons,
further,
electrons,
nucleus,
inner electron shells,
electrons,
EASIER.
Hence ionisation energy decreases down group.
What happens to the ionisation energy ACROSS A PERIOD?
Ionisation energy increases - harder to remove electrons.
The same —– of electrons is FILLED UP/EMPTYING across a period.
Outer ——— are CLOSER TO/FURTHER FROM the nucleus as they are more ——— to the increased ——– —— of the nucleus.
The outer ——— are EASIER or HARDER to remove.
Hence…
shell,
FILLED UP.
electrons,
CLOSER TO,
attracted,
positive charge.
electrons,
HARDER,
Hence ionisation energy increases.
Why are second ionisation energies always higher than the first?
Although electrons are being removed, the number of protons in the nucleus remains the same.
Therefore the pull on the remaining electrons is increased (making them harder to remove, so it takes more energy to remove a second electron than the first).
In energy terms, it’s not possible for atoms of group 2 elements to form a 3+ ion. Why?
It would mean removing a third electron from an energy level closer to the nucleus, which requires (too much) more energy.
The higher the ELECTRONEGATIVITY…?
The stronger the attraction of the atom for the shared electrons in a covalent bond.
Electronegativity values increase across a period. True or False?
Explain why. [4 marks]
True.
Same shell of electrons is being filled up [1]
There’s an increase in positive nuclear charge, pulling in electron shells [1]
Bonding electrons are closer to the nucleus. [1]
Causing electronegativity value to increase. [1]
Do the electronegativity values increase down a group?
Explain why. [3 marks]
No, electronegativity decreases.
Bonding electrons in outer shell are further away from the positive nucleus [1] and are screened from the full effect of the nucleus by the inner shells of electrons [1].
Causes electronegativity value to decrease. [1]
What does the term “Periodicity” (with reference to the Periodic Table) mean?
Repeating patterns/trends in the properties of elements.
Which type of bonding holds aluminium atoms together in a sample of the element?
Metallic
Why does fluorine have a smaller covalent radius than lithium?
Fluorine has more protons inside its nucleus.
*Covalent radius decreases moving across a period due to increasing nuclear charge pulling the outer electrons more strongly.
Why is there such a large increase between the second and third ionisation energies of calcium?
The third electron must be removed from a full energy level (requiring lots of energy).
Which of the following atoms has least attraction for bonding electrons? [1 mark]
A: Carbon
B: Nitrogen
C: Phosphorus
D: Silicon
D - Silicon
*Refer to Page 12 of the Higher Chemistry Data Booklet (Publication Date: 2021) for the “Electronegativities of Selected Elements”.
*Silicon has the lowest electronegativity value of the four.
The first 3 ionisation energies (kJ/mol) of aluminium are as follows:
1st: 578
2nd: 1817
3rd: 2745
Using this info, what’s the enthalpy change, in kJ/mol, for the following reaction? [1 mark]
Al⁺(g) –> Al³⁺(g) + 2e⁻
A: 1817
B: 2395
C: 4562
D: 5140
C: 4562
*Steps to work it out:
(1) Al⁺(g) –> Al²⁺(g) + e⁻
ΔH=1817kJ/mol
(2) Al²⁺(g) –> Al³⁺(g) + e⁻
ΔH=2745kJ/mol
Add (1) + (2)
= Al²⁺(g) appears on both sides of ionisation reaction, so they cancel out.
Leaving us with…
Al⁺(g) –> Al³⁺(g) + 2e⁻
ΔH=4562kJ/mol
(units can also be written as kJmol⁻¹)
Explain why the COVALENT RADIUS of sulfur is smaller than that of phosphorus. [1 mark]
Sulfur has more protons in its nucleus.
OR
There is increased nuclear attraction for electrons.
OR
Increased nuclear charge
The covalent radius is a measure of the size of an atom.
Explain why covalent radius decreases across period 3 - from sodium to chlorine. [1 mark]
(The electron shells are pulled closer because) nuclear charge increases/the number of protons in the nucleus increases.
Explain FULLY why the covalent radius of sodium is larger than the ionic radius of sodium. [2 marks]
TWO POINTS ARE REQUIRED.
Sodium atom loses an electron to become sodium ion [1]
AND
Sodium ion only has two electron shells (2,8) whereas the sodium atom has three electron shells (2,8,1). [1]
Explain why the first ionisation energy decreases going down Group 1. [1 mark]
Outer electrons less strongly attracted to the nucleus
OR
Outer electrons are more shielded from the nuclear pull as atom increases in size.
Explain FULLY why the second ionisation energy is much greater than the first ionisation energy for group 1 elements. [2 marks]
The electron removed is being removed from a full outer shell or from an electron shell closer to the nucleus. [1]
The electron being removed by the 2nd ionisation energy is more strongly attracted to the nucleus (less shielded). [1]
*The electron being removed has to be removed from a full electron shell [1], which is closer (more strongly attracted) to the nucleus [1].
Explain why the first ionisation energy increases across period 3 of the periodic table. [1 mark]
Increasing number of protons
OR
Increasing nuclear charge
Write an equation, including state symbols, for the SECOND ionisation energy of magnesium. [1 mark]
Mg⁺(g) –> Mg²⁺(g) + e⁻
*Must include state symbols
State what is meant by the term ELECTRONEGATIVITY. [1 mark]
The measure of
attraction a nucleus has for the electrons in a bond.
Explain why electronegativity values decrease going down group 7 (halogens) of the periodic table. [1 mark]
(More shells) so increased shielding.
OR
Covalent radius increases/atom size increases/more shells so attraction of the nucleus/protons for the OUTER ELECTRONS decreases.
Explain the decrease in atom size going across period 3 of the periodic table - from sodium to argon. [1 mark]
Increasing number of protons (in the nucleus).
OR
Increasing nuclear charge.
Explain why the first ionisation energy increases going from lithium to neon. [1 mark]
Increasing number of protons
OR
Increasing nuclear charge
Explain why the first ionisation energy of potassium is less than the first ionisation energy of lithium. [1 mark]
(More shells) so increased shielding
OR
Covalent radius increases/atom size increases/more shells so attraction of the nucleus/protons for the OUTER ELECTRONS decreases.
Write an equation for the second ionisation energy of nitrogen. [1 mark]
N⁺(g) –> N²⁺(g) + e⁻
*Must show state symbols.
Explain FULLY the increase between the 5th and 6th ionisation energies of nitrogen. [2 marks]
The 6th electron being removed has to be removed from a full electron shell, which is closer and more strongly attracted [1] to the nucleus [1].
USING YOUR KNOWLEDGE OF CHEMISTRY, comment on similarities and differences in the patterns of physical and chemical properties of elements in both Group 1 and Group 4 of the Periodic Table. [3 marks]
Include descriptions/elaborate on:
Types of Bonding,
Types of Structures,
Melting & Boiling Points,
Covalent Radii,
Ionisation energies,
Electronegativities, etc.
You can find a helpful breakdown of an example answer from this link -> https://youtu.be/O-LTNmZyRm0
The melting point of NON-METAL elements depends on structure and bonding.
USING YOUR KNOWLEDGE OF CHEMISTRY, comment on this statement.
Mention:
ALL COVALENT BONDING TYPES -
(1) Monatomic (Group 0) elements/(gaseous) structures
(2) Covalent Molecular (diatomic) elements/structures and the P₄, S₈, Fullerenes (eg. C₆₀) structures
(3) Covalent (diamond, graphite, silicon, boron) Network bonding and structures
*ESSENTIAL THAT YOU MENTION INTERMOLECULAR (Weak London Dispersion) FORCES for all 3 bonding types.
More electrons? Stronger London Dispersion Force.
A helpful breakdown of the question can be found here -> https://youtu.be/C0uHTjgiJvQ
An element contains covalent bonding and London dispersion (intermolecular) forces.
The element could be:
A: Boron
B: Neon
C: Sodium
D: Sulfur
[1 mark]
D - Sulfur
*Sulfur is a small individual molecule with strong London dispersion (intermolecular) forces.
Explain the decrease in covalent radius in period 2 of the Periodic Table - from Nitrogen to Fluorine. [1 mark]
(The electron shells are pulled closer because) nuclear charge increases/the number of protons in the nucleus increases.
First ionisation energies decrease going down a group.
State what is meant by the term FIRST IONISATION ENERGY. [1 mark]
The energy required to remove one mole of electrons from one mole of gaseous atoms.
Explain why the first ionisation energy of the group 7 elements decrease going down the group. [1 mark]
Outer electrons less strongly attracted to the nucleus
OR
Outer electrons are more shielded from the nuclear pull as atom increases in size.
The boiling points (*C) of three hydrogen “halides” are as follows:
Hydrogen chloride = -85
Hydrogen bromide = -66
Hydrogen iodide = -35
Explain FULLY why the boiling point increases from hydrogen chloride to hydrogen iodide. [2 marks]
The London dispersion (intermolecular) forces become stronger/increase (in moving from HCl to HI).
The number of electrons in the molecules increases (from HCl to HI).
Atoms of different elements have different attractions for bonding electrons.
Electronegativity is a measure of the attraction an atom involved in a bond has for the electrons in the bond.
USING YOUR KNOWLEDGE OF CHEMISTRY, discuss the importance of electronegativity in bonding, structure and properties of compounds. [3 marks]
Talk about electronegativity value differences for:
- Non-Polar Bonds
- Polar Bonds
Mention mp/bps and van der Waals Forces!
Bonus points, if there are diagrams.
Compare the ionic sizes of Mg²⁺ and S²⁻. Which ion is bigger?
Steps:
(1) Write down the electron arrangement of Mg. Now minus 2 electrons (because it’s Mg²⁺). The electron arrangement is now 2,8. –> 2 electron shells
(2) Write down the electron arrangement of S. now add 2 electrons (because it’s S²⁻).
The electron arrangement is now 2,8,8. –> 3 electron shells
(3) Think about covalent radii. The number of protons stays the same.
Mg²⁺ = stronger nuclear charge = smaller ion size.
S²⁻ = weaker nuclear charge because there’s more distance between outer electrons and nucleus = larger ion size.
Explain the trend in atomic size
a) going down a group
b) going across a period
a) Atomic size increases going down a group due to the increasing number of electron shells.
b) Atomic size decreases going across a period due to the increasing number of protons/nuclear charge pulling the electrons in (decreasing the atomic size).
Why is the Mg ion smaller than the Mg atom?
Mg ion is smaller than Mg atom because the Mg atom contains 1 more electron shell than the ion.
Is the fluorine atom smaller than the fluorine ion? Are they the same size?
Fluorine atom is bigger than the fluorine ion because Fluorine ion is subject to increased nuclear charge.
Explain why Na is bigger than Cl.
Cl has more protons (thus stronger nuclear charge) than Na. Hence Cl has a smaller atomic size.
Explain why Na+ is smaller than Cl-.
Na+ is smaller than Cl- because it has one less occupied energy level.
Explain why Rb is bigger than K.
Rb has more occupied energy levels than K.
Explain why a phosphorus ion is bigger than a silicon ion.
Silicon ions have a larger nuclear charge (less shielding - electrons get more pulled in).
Calculate the energy required for Al (g) –> Al3+ (g)
2745kJ/mol
*You can just take the value straight from the databook as it’s going from an atom to an ion.
Calculate the energy required for K+ (g) –> K3+ (g)
4420 - 419 = 4001kJ/mol
*Take the first ionisation energy value away from the third ionisation energy value in this case… :)
The greater the number of occupied energy shells, the ——- the covalent radius.
bigger
Some Periodic Tables show hydrogen at the top of Group 1 and/or Group 7.
Suggest why some chemists think it appropriate to place hydrogen:
a) at the top of Group 1
b) at the top of Group 7
a) Hydrogen can be placed at the top of Group 1 as it has an electron arrangement similar to the Alkali Metals.
b) Hydrogen can be placed at the top of Group 7 as it only needs one electron to fill its outer shell like the halogens.
The electronegativities of elements are listed on page 12 of the data book. Why is there no value for the noble gases?
The noble gases do not form to make compounds as all of their electrons are bonded. So no electronegativity can be measured.
