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What is the modern periodic table based on?
Dimitri Mendeleev's periodic table
How did Dimitri Mendeleev order the elements?
By their atomic masses
How did Dmitri Mendeleev line up elements in groups?
By their similar chemical properties
What were remarkable predictions of Mendeleev?
He predicted elements that hadn't been discovered
As of 2014, how many elements are in the periodic table?
114 elements
How many vertical groups are there in the periodic table?
18 groups
How many periods are there in the periodic table?
7 Periods
How are elements arranged on the periodic table today?
In order of increasing atomic numbers (number of protons)
What is each horizontal row of the periodic table called?
A period.
What is each vertical column of the periodic table called?
A group
What does the position of elements on the periodic table prove?
Proves the physical and chemical properties of that element
What is the trend across a period in the periodic table?
- As you move across the period, number of protons increases
- Metals change to non-metals
What is periodicity?
Repeating trend of properties across a period
What is the trend across period 2?
2s sub-shells filled with two electrons followed by a 2p subshell
What is the periodicity trend down a group?
- All elements have the same number of electrons in the outer shell
- All elements have the same type of orbitals on the outside
What are four properties that depend on periodicity?
- Electron configuration
- Ionisation energy
- Structure
- Melting points
What is ionisation energy of an atom?
The energy needed to remove an electron from an atoms' outer shell
What does ionisation energy measure?
The energy required for atoms to lose electrons to form positive ions
What is the definition of the first ionisation energy?
Energy required to remove one mole of electrons from each atom in one mole of GASEOUS atoms of an element to form one mole of gaseous 1+ ions
What are the factors affecting ionisation energy?
- Atomic radius
- Nuclear charge
- Electron shielding/screening
Why is energy required to form an ion?
Energy must be supplied to an electron to overcome attraction from the positive nucleus
How does increasing atomic radii affect the ionisation energy?
- The larger the atomic radii, the lesser the effect the positive nucleus has on attracting negative electrons
- Increased atomic radii therefore decreases ionisation energy
How does increasing nuclear charge affect the ionisation energy?
- The higher the nuclear charge the larger the positive attractive force has on negative electrons
- This therefore increases ionisation energy as more energy is required to overcome the forces of attraction
What exactly is electron shielding?
Repelling effect towards the outer shell from multiple negative electrons in the inner shell
How does electron shielding/screening affect the ionisation energy?
- Inner shells of electrons repel the outer shell electrons.
- As you increase the number of inner shells you increase a larger shielding effect, reducing the effect of nuclear attraction
- This decreases ionisation energy
What are successive ionisation energies?
A measure of the energy required to remove each electron
What three reasons are there for why a new successive ionisation energy is greater than the previous successive ionisation energy?
- As you remove each electron, less repulsion remains between the electrons and inner electrons (decreased electron shielding), this increases ionisation energy
- Additionally, as you remove each electron, positive nuclear charge will therefore distribute more positive attraction to the remaining electrons, increasing nuclear effect, also increasing ionisation energy
- This also decreases atomic radii of the atom causing even further attraction, increasing ionisation energy
What is the definition of the second ionisation energy?
Energy required to remove one mole of electrons from each ion in one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions
How is a first ionisation energy displayed? Use helium as the example
He (g) --> He+(g) + e-
How is a second ionisation energy displayed? Use helium as the example
He+(g) --> He2+ (g) +e-
What can successive ionisation energies allow in terms of predictions of the periodic table?
- Number of electrons in the outer shell
- Group of the element in the periodic table
- Identity of the element
What three factors explains the increase or decrease of first ionisation energies between elements?
- Atomic radii
- Electron shielding
- Nuclear charge
What are factors of ionisation energies that can cause dips? (Name the obvious factor and the subtle factor)
- The jump between periods
- The changes in electron sub-shells
- Certain removal of electrons from the orbitals
What is the pattern in first ionisation energy for the first 20 elements as you increase the element?
- A general increase in the first ionisation energy as you go across the period
- A sharp decrease in first ionisation energy between a jump of periods
What is the trend as you increase the number of inner shells down groups? (talk about three important ionisation energy factors)
- Atomic radius increases as you go down the group
- More inner shells occur so electron shielding also increases
- Nuclear charge increases but is outweighed by atomic radii and electron shielding
- Therefore nuclear attraction on outer electrons decreases
- Therefore first ionisation energy decreases
What is the trend as you increase atomic number?
- Nuclear charge and attraction increases
- Atomic radii decreases
- Electron shielding does increase however is countered by atomic radii and nuclear charge
- First ionisation therefore increases
What happens to the first ionisation energies as sub shells change across the right of a period?
First ionisation decreases
Why is it easier to remove one electron from one electron filled p orbital than one electron from a full s orbital? What happens to the first ionisation energy?
- P orbital has a higher energy level than an S orbital
- P orbital is further away from the nuclei
- First ionisation energy decreases
Why is it easier to remove one electron from an orbital with an electron pair than in an orbital of the same type of subshell with no electron pairs? What occurs to the first ionisation energy?
- Electron pair repulsion makes it easier to remove an electron from an electron pair in an orbital
- First ionisation energy deceases
What metal is not a solid at room temperature?
Mercury
What bonding does atoms of the same metal element form?
Metallic bonding
What are the negative subatomic particles in metallic bonding?
Delocalised electrons.
What are the positive ions in metallic bonding? How are they arranged as part of their structure?
- Metals are positive cations,
- Cations occupy fixed positions in the lattice
How should the structure (arrangement) of metals be described?
- Arrangement of fixed positive cations in a sea of delocalised electrons that act as mobile charge carriers
- This arrangement forms a metallic lattice structure
Is it possible to tell which electron originated from a specific metal ion?
No, impossible
Must the charges in metallic bonding balance?
Yes
What happens when electricity is applied to a metal? Why?
- Metal conducts the electricity
- When a voltage is applied to a metal, delocalised electrons can move through the structure and act as mobile charge carriers
Do metal, like ions, need to be in a liquid state to carry charge?
No, the metals can be in a solid state
What are the four properties of metals?
- High melting and boiling points
- High thermal conductivity
- High electrical conductivity
- Malleability and ductility
Why do most metals have high melting and boiling points?
Lots of energy is required to overcome the very strong electrostatic attraction between the positive cations and the sea of delocalised electrons. Therefore high melting point is present
Why do metals not dissolve?
Any interactions between metals and covalent bonds would result in a reaction forming ions
What is malleability?
Metals can be beaten into shapes through pressure
What is ductility
Metals that can be drawn out or stretched
Why are metals ductile and malleable?
- Delocalised electrons allow atoms or layers to slide over each other
- This gives the property of malleability and ductility
What elements have a giant metallic structure?
ALL METALS INCLUDING
- Lithium
- Beryllium
- Sodium
- Magnesium
- Aluminium
What elements form giant covalent structures?
- Boron
- Carbon
- Silicon
What occurs to the melting points of elements as you across the period?
- Melting points and boiling points increases in metals as you go across
- However if giant covalent forces present then melting points and boiling points also increases
How does diamond form?
One carbon atom is covalently bonded to three other carbon atoms
What happens to the melting points as you change structures from giant metallic to giant covalent?
Melting point and boiling points increases substantially
What happens to the melting points as you change the structures from giant covalent to simple covalent?
Melting points and boiling points decrease
Why can't diamond conduct electricity?
Diamond cannot conduct electricity because there are no mobile electrons to carry the electrical charge
Why can graphene and graphite conduct electricity?
Delocalised electrons which act as mobile charge carriers
What are group 2 metals also known as?
Alkaline earth metals
Which group 2 metal is the least reactive?
Beryllium
Which group 2 metal is the most reactive?
Barium
Are group 2 metals oxidising or reducing agents?
Reducing agents
What ion will every group 2 metal form?
2+ cations
When group 2 metals react to form 2+ ions what will it lose to another species?
2 electrons
How does a group 2 metal react with oxygen?
2Mg + O2 --> 2MgO
How does a group 2 metal react with water?
Mg + 2H20 --> Mg(OH)2 + H2
What is the trend in reactivity in group 2
Reactivity increases as you go down group 2
Can beryllium react with water?
No, least reactive
How does a group 2 metal react with a dilute acid?
Mg(s) + 2HCl(aq) --> MgCl2(aq) + H2(g)
What are the physical properties of group 2 metals?
- High melting points and boiling points
- Light metals with low densities
- Form white colourless compounds
Why does reactivity increase as you go down group 2?
As it becomes easier to remove electrons from the s sub shell
Why does it become easier to remove an electron from the s-subshell as you go down group 2?
- As atomic radii increases
- Electron shielding increases
- Nuclear charge increases HOWEVER THIS IS OUTWEIGHED BY INCREASED ATOMIC RADII
(MUST GET THIS RIGHT)
How many successive ionisation energies does a group 2 element undergo?
2 successive ionisation energies
Describe the first and second successive ionisation energies of a group 2 element. Use calcium as an example.
Ca(g) --> Ca+(g) + e- (first ionisation energy)
Ca+(g) --> Ca2+(g) + e- (second ionisation energy)
When a metal and water react together, what does the main product (with the metal) dissociate?
OH- ions
As you go down group 2, does the solubility increase or decrease in group 2 metal hydroxides?
Solubility increases
(Remember, metals are not soluble however metal oxides are soluble)
How do you test the difference of solubility of metal hydroxides as you go down the group?
1) Add a spatula of a group 2 oxide to water
2) Shake the mixture being careful not to dissolve all metal hydroxides
3) Measure the pH and compare between different metal oxides
What is the least soluble group 2 metal hydroxide?
Magnesium hydroxide
(Beryllium cannot react with water therefore It can't form a hydroxide)
What group 2 hydroxide compounds is used in agriculture?
Calcium hydroxide known as lime
How does calcium hydroxide (lime) decrease acidity of soil?
Calcium hydroxide neutralises acid in the soil and adds OH- ions
Why can group 2 metal hydroxides be used for indigestion (or acid reflux)?
Group 2 bases dissociates OH- ions allowing neutralisation of the stomach's pH
What are possible group 2 metal hydroxide examples that can be used for neutralisation of stomach pH
- Magnesium hydroxide
- Calcium hydroxide
What is the most reactive non-metallic group on the periodic table?
Group 17; The halogens
What common compounds can halogens exist as on earth? (think type of compound)
Ionic compounds, i.e. NaCl
At RTP what state and what bonding do halogens exist as? (think their covalent structure)
Diatomic gas molecules;
- i.e. iodine exists as I2
- Chloride exists as Cl2
What are the relative melting and boiling points of halogenic diatomic molecules?
Diatomic halogens have relatively low melting and boiling points
As you go down the group, does the melting points of halogens increase or decrease?
Melting and boiling points increase
Why is the trend of melting and boiling points of halogens increasing as you go down the group?
- More electrons leads to stronger and more occurrences of London forces in between molecules
- More energy is therefore required to overcome these intermolecular forces causing the melting point and boiling point to increase
Does the reactivity and oxidising power decrease or increase as you go down the group?
Reactivity and oxidising ability decreases as you move down the group
Why does reactivity and oxidising ability decrease as you move down the group?
- Atomic radii increases distance and decreases nuclear attraction
- Electron shielding decreases nuclear attraction
- These two factors outweigh increase in nuclear charge
- Ability to gain an electron in the p-sub shell and form -1 ions decreases due to the lower nuclear attraction
How do displacement reactions of halogen ions occur?
A more reactivate halogenic diatomic molecule will oxidise and displace a halide ion of less reactivity
What happens when chlorine and water react?
Pale green colour emerges
What happens when chlorine and cyclohexane reacts?
Pale green colour emerges
What happens when bromine and water reacts?
Orange colour emerges
What happens when bromine and cyclohexane reacts?
Orange colour emerges
What happens when iodine and water reacts?
Brown colour emerges
