3.1 - The periodic table and periodicity Flashcards
Define the term ‘periodicity’ ?
Periodicity - The study of properties of elements across each period that show a repeating pattern
Explain how elements are arranged in the periodic table ?
- They are arranged in increasing atomic number
- They are arranged in periods
- They are arranged in groups
What does the arrangement of elements in periods represent ?
- It represents the number of shells atom contains
- It represents repeating trends in physical and chemical properties
What does the arrangement of elements in groups represent ?
- It represents the number of out shell electrons an atom contains
- It shows elements that have similar chemical properties
Define the term ‘atomic radius’ ?
Atomic radius - Half the distance between the nuclei of two adjacent atoms
What is the trend in atomic radius across a period ?
As you go across the period, the atomic radius decreases
Explain the trend in atomic radius across a period ?
- As you go across the period, the number of protons increases but the shielding remains the same
- This means that the effective nuclear charge increases which means the outer shell electrons are more attracted to the positive nucleus
- This means they are pulled more towards the nucleus causing the atomic radius to decrease
Identify the s, p, d and f blocks on the periodic table ?
How do you determine what block is assigned to an element ?
- An element can be assigned to the s, p, d or f block according to where it highest energy electrons are
- Eg. Magnesium is in the s-block since its outer electrons / highest energy electron is in the 3s orbital ( 1s2 2s2 2p6 3s2 )
Explain the increasing trend in melting points for the first three elements across periods 2 and 3 ?
- They have a giant metallic lattice structure
- As you go across the period, the number of outer shell electrons increases
- This means metals form positive metal ions with a greater positive charge
- This means the positive metal ions are more attracted to delocalised electrons which means metallic bonding becomes stronger
- This means more energy is required to overcome strong metallic bonds which causes melting/ boiling point to increase
Explain the trend in melting points for the fourth element across periods 2 and 3 ?
- They have a giant covalent lattice structure
- Therefore, atoms are only held together by many strong covalent bonds which require a lot of energy to overcome
- This means they have a very high melting/ boiling point
Explain sudden decrease between the fifth and seventh element across periods 2 and 3 ?
- They have a simple molecular lattice structure
- Although atoms are held together by strong covalent bonds, molecules are held together by weak London forces which don’t require a lot of energy to overcome
- This means they have a very low melting/ boiling point
Explain the increasing trend in melting points between the fifth and seventh element across period 2 ?
- As you go across the period, the number of electrons increases
- This results in larger dipoles and therefore stronger London forces which require more energy to overcome
- Therefore, the melting/ boiling point increases across the period
Explain the increasing trend in melting points between the fifth and seventh element across period 3 ?
- As the size of the molecule increases, the number of electrons present increases
- This results in larger dipoles and therefore stronger London forces that require more energy to overcome
- Therefore larger molecules have higher melting/ boiling points
Explain the trend in melting points for the eight element across periods 2 and 3 ?
- These elements are monoatomic
- This means they have a low number of electrons which results in very small dipoles
- This means they have very weak intermolecular forces between atoms which don’t require a lot of energy to overcome resulting in very low melting/ boiling points
What is the melting point determined by ?
- It is dependant on the size of the molecule
- As the size of the molecule increases, the number of electrons present increases
- This results in greater dipoles and therefore stronger London forces which require more energy to overcome
- This causes the melting/ boiling point to increase
What is the structure of metals ?
- A Giant metallic lattice structure
- A 3D arrangement of positive metal ions surrounded by a sea of delocalised electrons held together by strong metallic bonds
Define the term ‘metallic bond’ ?
Metallic bond - The strong electrostatic attraction between positive metal ions and delocalised electrons
What are the features of metals ?
- High melting/ boiling point
- Electrical and thermal conductors
- Insoluble
Explain why metals have high melting/ boiling points ?
- Metals contain strong metallic bonds that require a lot of energy to break/ overcome
- This means they have a large melting/ boiling point
Explain why metals are electrical conductors ?
- Metals contain delocalised electrons which can move
- This means metals can conduct electricity in the liquid and solid state
Explain why metals are insoluble ?
Metals do not dissolve
What is the structure of a simple molecule ?
- They have a simple molecular lattice structure
- Atoms are held together by strong covalent bonds while molecules are held together by weak London forces/ induced dipole-dipole interactions
What are the features of simple molecular substances ?
- Relatively low melting/ boiling point
- Cannot conduct electricity
Explain why simple molecular substances have a relatively low melting/ boiling point ?
- There are only strong covalent bonds between atoms while there are weak London forces between molecules
- London forces are less strong/ weaker than strong covalent bonds so less energy is required to overcome them
- This results in a low melting/ boiling point
Explain why simple molecular substances cannot conduct electricity ?
- Simple molecular substances do no contain moving electrons or mobile ions
- Therefore, they cannot conduct electricity
What is the structure of an ionic compound ?
- Ionic compounds have a giant ionic lattice structure
- It is a 3D arrangement of oppositely charged ions held together by strong covalent bonds
Define the term ‘ionic bond’ ?
Ionic bond - The strong electrostatic attraction between oppositely charged ions
What are the features of an ionic compound ?
- They have a high melting/ boiling point
- They are soluble
- They cannot conduct electricity in the solid state
- They can conduct electricity in the molten/ aqueous state
Explain why ionic compounds have a high melting/ boiling point ?
- Oppositely charged ions are held together by strong ionic bonds that require a lot of energy to overcome
- This means they have high melting / boiling point
Explain why ionic compounds are soluble ?
Ionic compounds dissolve in most polar solvents
Explain why ionic compounds are both poor and good electrical conductors ?
- In the solid state, ions are fixed in the giant ionic lattice structure therefore no mobile charge carriers and cannot conduct electricity
- In the molten/ aqueous state, ions are mobile so can conduct electricity
Define the term ‘covalent bond’ ?
Covalent bond - The shared pair of electrons
What is the structure of Giant covalent substances ?
- They have a giant covalent lattice structure
- A 3D network of atoms held together by strong covalent bonds
Name some Giant covalent substances ?
- Graphene
- Graphite
- Diamond
- Silicon
What is the structure of diamond ?
- Each carbon atom is covalently bonded to 4 other carbon atoms
- It is very hard, very strong and has a high melting point due to carbon atoms being held together by strong covalent bonds that take a lot of energy to overcome
What is the structure of graphite ?
- Each carbon is covalently bonded to 3 other carbon atoms
- One outer shell electron from each carbon atom becomes delocalised
- Each layer of graphite is held to the next by weak intermolecular forces ( slippery )
What is the structure of graphene ?
- Each carbon atom is covalently bonded to 3 other carbon atoms
- One outer shell electron from each carbon atom becomes delocalised, so it conducts electricity
- It is only one atom thick ( transparent + flexible )
What are the features of Giant covalent substances ?
- High melting/ boiling point
- Insoluble
- Not electrical conductors ( apart from graphene/ graphite )
Explain why Giant covalent substances have a high melting/ boiling point ?
- Atoms are held together by many strong covalent bonds that require a lot of energy to overcome
- This means they have a high melting/ boiling point
Explain why Giant covalent substances insoluble ?
- They do not dissolve in almost all solvents
- Arms are held together by strong covalent bonds which require too much energy to overcome to be broken by interacting with solvents
Explain why graphene and graphite electrical conductors ?
- In graphene/ graphite, each carbon atom donates one delocalised electron to the ‘sea’ of delocalised electrons
- Graphene/ graphite contain delocalised electrons that can move so can conduct electricity
What are the factors that affect electronegativity ?
- Nuclear charge
- Atomic radius
- Shielding
Define the term ‘electronegativity’ ?
Electronegativity - The power of an atom to attract the electrons in a covalent bond
Explain how nuclear charge affects electronegativity ?
- The more protons in the nucleus, the stronger the electrostatic attraction between the nucleus and bonding pairs of electrons
Explain how atomic radius affects electronegativity ?
- The smaller the atomic radius, the smaller the distance between the outer shell electrons and nucleus
- This means the electrostatic attraction between the nucleus and bonding pairs of electrons is stronger
Explain how shielding affects electronegativity ?
- As the number of shields increases there are more electrons between the nucleus and the bonding pairs of electrons
- This means the electrostatic attraction between the nucleus and bonding pairs of electrons in weaker electrons
What is the trend in electronegativity down the group ?
As you go down the group, the electronegativity decreases
Explain the trend in electronegativity down the group ?
- As you go down the group the atomic radius and shielding increases
- This means there is more distance and a larger number of electrons between the nucleus and the bonding pairs of electrons
- This weak rand the electrostatic force between the nucleus and bonding pairs of electrons
What is the trend in electronegativity across the period ?
As you go across the period, the electronegativity increases
Explain the trend in electronegativity across the period ?
- As you go across the period the atomic radius decreases and the number of protons increases and shielding remains the same meaning the effective nuclear charge increases
- This means the atomic radius decreases
- This means there is shorter distance between the nucleus and bonding pairs of electrons
- This means that the electrostatic attraction between the nucleus and bonding pairs of electrons is stronger
Define the term ‘first ionisation energy’ ?
First ionisation energy - The amount of energy required to remove one mole of electrons from one mole of atoms in the gaseous state at STP ( Standard conditions of temperature and pressure )
Write an equation for the first ionisation energy of Sodium/ Na ?
Na (g) –> Na+ (g) + e-
Do you know how to write an equation for the ionisation energies of different elements ?
Yes
What factors affect first ionisation energy ?
- Atomic radius
- Nuclear charge
- Shielding
How does atomic radius affect ionisation energy ?
- As the atomic radius increases, the greater the distance between the outer shell electrons and positive nucleus
- This means the outer shell electrons are less targeted to the nucleus so less energy is required to remove them
- This causes the ionisation energy to decrease
How does effective nuclear charge affect ionisation energy ?
- As the number of protons increase, the nuclear charge increases
- This can cause the effective nuclear charge to increase which means the outer shell electron is more attracted to the nucleus
- Therefore, more energy is required to remove the outer shell electron causing the ionisation energy to increase
How does shielding affect ionisation energy ?
- As the shielding increases, the number of shells increases
- This means the outer shell electrons are less attracted to the nucleus, meaning less energy is required to remove the outer shell electron
- This causes the ionisation energy to decrease
What is the trend in ionisation energy down the group ?
As you go down the group, the ionisation energy decreases
Explain the trend in ionisation energy down the group ?
- As you go down the group, the atomic size increases and the shielding increases
- This means that the outer shell electrons and further away from the nucleus and therefore less attracted to the nucleus
- Therefore, less energy is required to remove them causing the ionisation energy to decrease
What is the trend of ionisation energies across a period ?
As you go across a period, the ionisation energy generally increases
Explain the trend in ionisation energy across a period ?
- As you go across a period, the number of protons increases but the shielding remains the same
- This means the effective nuclear charge increases which means the outer shell electrons are more attracted to the positive nucleus
- this means more energy is required to remove the outer shell electrons which causes the ionsaition energy to increase
Draw a graph showing the successive ionisation energies of an element ( ie. Al ) ?
Can you draw a graph showing the successive ionisation energy of an element ?
Yes
Interpret this graph, explain what element in the period two this graph represents ?
- There is a ‘jump’ in ionisation energy between the third and fourth electrons removed
- This means that the fourth electron is removed from an inner shell which means that the element has three outer shell electrons
- This means that the element is Boron since it is in group 3, period 3
Can you interpret values of successive ionisation energy and a successive ionisation energy graph ?
Yes
Draw a diagram representing the ionisation energies across a period ( eg. period 2 ) ?
Can you draw a graph representing ionisation energy across a period ?
Yes
Explain the anomaly in ionisation energies across the period for ie. Boron ?
- Boron has a lower ionisation energy than beryllium since it contains a 2p1 ( different shell for different periods ) electron which is slightly higher in energy
- This means less energy is required to remove it meaning it has a lower ionisation energy
Explain the anomaly in ionisation energies across the period for ie. Oxygen ?
- Oxygen has a lower ionisation energy than nitrogen since it contains four electrons in the p sub-shell
- This means that two electrons occupy on of the p orbital and repel due to like charges
- This means less energy is required to remove the 2p4 electron meaning it has a lower ionisation energy
