Module 3.1 Flashcards
Give an outline of the periodic table
- Consists of rows (periods- horizontally across)
- Consists of groups
- Arranged by increasing atomic number
In the periodic table, elements are arranged by …
Increase atomic number
What groups are fond in the s block?
Groups 1 and 2
What groups are found in the p block?
Groups 3-8
What are the names of the elements found in block d?
Transition metals
Define the term periodicity
Trends that occur in physical and chemical properties as we move across the periods of the periodict table
When talking about periodicity, list 5 different trends that we talk about
- Ionisation energy
- Melting Points/Boiling Points
- Reactivity
- Atomic Radius
- Electronegativity
Do all the elements within a group have similar reactions, why?
Elements in the same group have the same number of electrons on their outer shell. Therefore, they have similar chemical properties.
What was Döbereiner’s theory about?
Triads
What was John Newlands’ theory about?
Law of octaves
What was Mendeleev’s theory about?
Gaps, elements arranged in order of increase atomic mass
What is the modern day theory of the periodic table?
Elements are arranged in order of increasing atomic number
Explain Döbereiner’s theory behind his model of the periodic table
- Döbereiner grouped similar elements in TRIADS e.g. Li,Na and K.
- He ordered elements by atomic mass.
- The middle element in his triads had SIMILAR properties to the other TWO elements.
Explain John Newlands’ theory behind his model of the periodic table
- Newlands arranged elements in order of mass. He noticed that every 8th element was SIMILAR.
Explain Mendeleev’s theory behind his model of the periodic table
- He arranged all known elements by atomic mass
- He left GAPS, where no element fitted the repeating patterns
- Predicted the patterns of missing elements
Explain the theory behind the modern day periodic table ( Henry Moseley)
- Elements are arranged by increasing atomic number
- There are groups and periods
Define first ionisation energy
The amount of energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms.
Give the equation for the first ionisation energy of oxygen
O(g) → O^+ (g) + e^-
Give the equation for the first ionisation energy of magnesium
Mg(g) → Mg^+ (g) + e^-
Give the equation for the first ionisation energy of Sodium
Na(g) → Na^+ (g) + e^-
When talking about first ionisation energy what must you include?
- Charges
- Gaseous state
- The amount of energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms
The lower the ionisation energy, the easier/harder it is to form an ion?
Easier
What factors affect ionisation energy?
Nuclear charge
Atomic radius
Shielding
How does nuclear charge affect ionisation energy?
The higher the nuclear charge, the larger the electrostatic attraction between the nucleus and the outermost electrons. Therefore, more energy is required to remove 1 mole of electrons from 1 mole of gaseous atoms, so the as nuclear charge increase, so does IE.
Higher nuclear charge = Higher Ionisation energy
How does the atomic radius affect ionisation energy?
A increase in atomic radius means that the nuclear attraction between outermost electrons and the nucleus is smaller. Therefore, less energy is required to remove 1 mole of electrons from 1 mole of gaseous atoms, so as the atomic radius increases, ionisation energy falls.
Higher atomic radius = Lower Ionisation energy
How does the shielding affect ionisation energy?
As the number of electrons between the outermost electrons and the nucleus increases (As shielding increases), the outer electrons feel less attraction towards the nucleus. Therefore, less energy is required to remove 1 mole of electrons from 1 mole of gaseous atoms, so as shielding increases, ionisation energy falls.
More shielding = Lower ionisation energy.
Does Ionisation energy increase or decrease down the group, and why?
Down the group ionisation energy DECREASES. This is because elements down the group have extra electron shells compared to the ones above. The extra shells mean that the atomic radius is larger, so the outer electrons are further away (increasing shielding) from the nucleus, which greatly reduces their attraction to the nucleus.
Even though as you go down the group, the positive charge of the nucleus increases, why does ionisation energy still fall?
The effect caused by the positive charge is OVERRIDDEN by the effect of an increase in atomic radius and extra shielding
Does Ionisation energy increase or across the period, and why?
Across the period, ionisation energy INCREASES.
This is because across the period, the positive charge of the nucleus increases. This causes electrons to be pulled closer to the nucleus, making the atomic radius smaller. This means that the outer electrons are more strongly attracted to the nucleus. Therefore, more energy will be needed to ionise the element meaning across the period ionisation energy increases
Define the term second ionisation energy
The energy needed to remove an electron from each of one mole of 1+ ions in a gaseous state.
List reasons why ionisation ionisation energies higher than the one before?
- Less repulsion (Greater nuclear attraction (due to a fall in atomic radius))
- Positive nuclear charge outweighing the effect of the electron every time the an electron is removed
- Atomic radius decrease
Explain why successive ionisation ionisation energies higher than the one before?
- Less repulsion (remaining electrons drawn slightly closer to the nucleus)
- Positive nuclear charge outweighs the negative effect every time an electron is removed
- As the distance of each electron from the nucleus decreases slightly, the nuclear attraction increases. Therefore, more energy is needed to remove each successive electron
Define Metallic bonding
The strong electrostatic attraction between positive ions and delocalised electrons.
Give an overview of metallic bonding
- Sea of delocalised electrons
- Giant lattice of positive metal ions
- Strong electrostatic forces between positive ions and sea of free electrons.
List the properties of metallic bonding
- High melting point/High boiling point
- Can conduct electricity (when solid AND when molten
- Malleable
- Insoluble
Do metals have a high melting/boiling point or a low one? Why?
Compounds with metallic bonding have a HIGH melting/boiling point.
This is because the strong electrostatic attractions between the positive ions and the delocalised electrons require a lot of energy to overcome.
Can metals conduct electricity? Why?
Compounds with metallic bonding can conduct electricity when solid AND when molten. This is because the DELOCALISED ELECTRONS can move and carry charge.
Are metals malleable? Why?
Yes there are.
This is because as there are no bonds holding specific ions together, the metal ions can slide past each other when the structure is pulled.
Are metals soluble or insoluble?
Insoluble.
Metals are insoluble except in LIQUID METALS. This is because of the strength of metallic bonds.
What is the structure formed by metals?
Giant metallic lattice
What is meant by a giant metallic lattice?
A lattice of cations fixed in position surrounded by a ‘sea’ of delocalised electrons
What must you do when drawing metallic lattices?
- Balance the charges over the whole structure (e..g. if its a 2+ metal, have 2 delocalised electrons per 2+ ion)
- Draw at-least 3 rows
- Show charges on ions
- Show the delocalised electrons
Define covalent bonding
(insert)
What is the name of the structure that diamond, graphite, graphene and silicon form?
Giant COVALENT lattices
What is meant by a giant covalent lattice?
Huge networks of covalently bonded atoms
Outline the structure of diamond
A form of carbon atoms where each atom forms 4 other carbon atoms around it.
List the properties of diamond
- Very high melting point
- Can’t conduct electricity
- Very hard
- Insoluble
Explain why diamond has a very high melting point?
The extremely strong covalent bonds require a lot of energy to melt. As there are many strong covalent bonds in diamond, a lot (state the temperature, which will probably be given in the question) of energy is needed to melt the strcuture.
Explain why diamond can’t conduct electricity
All the outer electrons are held up in localised bonds
Outline the structure of graphite
Graphite is a form of carbon which forms a giant covalent lattice where carbon atoms are arranged in sheets.
List the properties of graphite
- Very high melting point
- Can’t conduct electricity
- Very soft
- Insoluble
Explain why graphite has a very high melting point
The extremely strong covalent bonds in the sheets require a lot of energy (state the temperature, which will probably be given in the question) to melt.