Unit 3. Periodic Table Flashcards
What is the nuclear charge
The number of protons in the nucleus of an atom. Basically the atomic number
Name of vertical columns and horizontal row
Group and period
What does the period number tell you?
The number of occupied main energy levels in an atom
What does the group number tell you?
The number of valence electrons in an atom of that element
Where are metals, non-metals and metalloids located on the periodic table?
Metals located mainly on the left and middle in the s-, f- and d-block.
Non-metal elements are found on the right side in the p-block.
Metalloids (semi-metals) are located in a diagonal staircase that forms a boundary between metals and non metals.
Name the name of the groups in the periodic table and where they are found, as well as their characteristic properties
Alkali Metals- Group 1- most reactive, react strongly with water
Alkaline earth metals- Group 2- reactive but less reactive then group 1
Halogens- Group 17- most reactive group of non-metals
Noble gases- Group 18- very unreactive monatomic gases
Transition metals- D-Block region (apart from zinc)- relatively stable with only moderate reactivity useful for construction
What are the lanthanides and actinides
They are metallic elements that make up the f-block. Their two rows are separated from the rest of the periodic table. Atomic numbers from 89-103, they are radioactive. Lanthanides (57-71) are knows as rare earth metals
What are the s- p- d- and f- block directly correlated to?
The electron configuration with the sub levels: s,p,d,f. The block in which an element is tell us which sub-level is in the process of being filled.
How do you find the v.e of group 13 to 18?
Substract ten from group number.
What is atomic radii
The size of an atom
Trend in atomic radii?
Period
Starting with sodium (11) it decreases across the period until it reaches chlorine (17). Bc we gain an electron and a proton as we go across a period, electron shielding remains more or less constant, as the electrons fill the same main energy level. Therefore there is an increased attraction between the nucleus and outer electrons which pulls them closer.
Group
Going down a group, each element’s v.e occupies another extra main energy level which increases the atomic radii.
What is electron shielding?
Electron shielding occurs when outer electrons are shielded from the attraction of the nucleus by inner electrons knows as shielding electrons.
Which are more likely to become anions and cations, metals or non metals?
Metallic elements tend to loose electrons and form cations.
Non metallic elements tend to gain electrons and form anions
What are isoelectronic species?
Isoelectronic species have the same electron configuration. Example: Na+, Mg2+. They have different numbers of proton but same of electrons.
What is the ionic radius
The distance between the nucleus of an ion and the outermost energy shell of the ion
When does the ionic radius increase?
When the number of proton decrease
More positively charged cations have the smallest ionic radius, and more negatively charged anions have the largest radius.
Explain ionic radius trend in period 3
- first 4 cations have lost their outer energy levels of elections and now occupy one less energy level. Since the nucleus is pulling on fewer electrons, the attraction increases between the nucleus and the v.e, which reduces the ionic radius.
- anions are bigger than their parent atom as they gain electrons which reduces the attraction by the nucleus. The extra electrons also increase repulsion between electrons. This all makes the ionic radii increase.
- across a period the ionic radii decreases as the nuclear charge increases (protons higher)
What is effective nuclear charge
The attraction felt by v.e which are further away from the nucleus and are shielded by other electrons.
How do you calculate the effective nuclear charge?
Take sodium (11). The 11th electron in the last 3s sub-level is shielded by the other 10 electrons, which are closer to the nucleus, from the 11 protons. Therefore do 11 protons minus 10 electrons. That v.e effective nuclear charge is 1+.
Nuclear charge trend
Period
- across increases as proton number increases but electron number in inner energy levels remains the same. It increases by one.
Group
- remains constant , as the attraction from the increasing numbers of protons in the nucleus is offset by the increase in occupied energy levels down the group.
What is first Ionisation energy?
Ionisation Energy is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions.
(A measure of the attraction between the positively charged nucleus and the negatively charged valence elections.)
General trend for first Ionisation energy
Increase across period bc of increase in nuclear charge and decrease in atomic radii, which results in an increased attraction between the nucleus and the v.e
Decrease down a group is bc of the increase in atomic radius which causes weaker attraction between the nucleus and v.e of an atom.
Exceptions to first ionization energy
Decrease in ionization energy from beryllium to boron, magnesium to aluminum, nitrogen to oxygen and phosphorus to sulfur.
Beryllium has the electron configuration 1s2 2s2
Boron has the electron configuration 1s2 2s2 2p1
Electrons in p orbitals are of higher energy and further from the nucleus than electrons in s orbitals, therefore they require less energy to remove. The same explanation can be applied for the drop in ionisation energy from magnesium to aluminium
Nitrogen has the electron configuration 1s2 2s2 2p3
Oxygen has the electron configuration 1s2 2s2 2p4
In oxygen, the electron is removed from a doubly occupied p-orbital. An electron in a doubly occupied orbital is repelled by the other electron and requires less energy to remove than an electron in a half-filled orbital. The same explanation applies for the decrease in ionisation energy from phosphorus to sulfur, except that the elect
What is electronegativity?
Electronegativity is the attraction of an atom for a bonding pair of electrons.
Trend of electronegativity
Increases across a period because of the increase in nuclear charge and the decrease in atomic radius.
Down a group electronegativity decreases bc of increase in atomic radius. The increase in nuclear charge down a group is counteracted by the increase in shielding caused by extra energy levels.
What is first electron affinity
The first electron affinity is the energy released when one mole of electrons is added to one mole of gaseous atoms to form one mole of gaseous 1- ions.
Are electron affinities exothermic or endothermic?
First Ionisation energies are exothermic bc energy is released when an electron is added to a neutral atom.
Second electron affinities are endothermic bc the repulsion that occurs when a second electron is added to an already negative ion.
Which elements have the highest electron affinities?
Group 17, they are the most exothermic as they are relatively small atoms that can accommodate an electron in their unfilled outer energy level.
Trend for electron affinity
Period
Across a period electron affinity increases, as we progress towards the filling of the outer valence shell of the atom. Group 17 releases more energy when aiming an electron as it achieves a filled valence shell and is more stable.
Group
Down the group the electron affinity decreases, as the additional electron gained is entering a new energy level further from the nucleus. This added electron has a weaker attraction to the nucleus and therefore releases less energy when added.
How are metals characterized?
- having fewer v.e
- larger atomic radii
- lower electronegativity and lower Ionisation energies
- tendency to loose v.e easily
How are non metallics characterized
- having more v.e
- smaller atomic radii
- higher electronegativity and Ionisation energy
- tend to gain electrons.
Trend of metallic character
Decreases across period
Increases down a group
Trend in melting point in period 3
Across a period the structure and bonding gradually changes from metallic to giant covalent to molecular covalent.
Sodium, mg and aluminum are metallic elements. the strength of the metallic bonding increases from sodium to aluminum, consequently so does the melting point.
The melting point reaches a peak with silicone (right after Aluminum), due to its giant covalent structure.
The elements to the right of silicone are non metallic, and although the covalent bonds within each molecule are strong, there are only relatively weak intermolecular forces between the molecules, so they have lower melting points.
What does the melting point rely on?
Intermolecular forces and structure
Trend of bonding and structure in period 3 oxides.
Period 3 oxides show a gradual decrease in metallic character across the period.
Metal oxides have giant ionic structures, which are sold under standard conditions due to the strong electrostatic attractions between the ions.
Non metals oxides are molecular compounds that exist as individual molecules. Bc of the weak intermolecular forces between their molecules, they are usually gas or liquids under standard conditions. They are giant covalent.
Na2O is ionic bc of the difference in electronegativity
But Cl2O is covalent bc both element are very electronegative an therefore the difference is less.
Acid base trend in period 3 oxides
All oxides before aluminum are basic. Al is amphoteric. Anything after Al is acidic.
Note: when term basic, amphoteric and acidic used, it means that they form basic, amphoteric and acidic solutions when reacted with water.
Trends of alkali metals (group 1)
- they are very soft
- first three elements have low densities that they can float on water
- melting points and boiling point are quite low down the group due to the metallic bonding getting weaker as the ionic radii of the metals increase
- they are stored in oil to prevent them from reacting with the air
- they easily loose their v.e due to low first ionization energy
- they undergo vigorous reactions with water to from metal hydroxide and hydrogen gas, which forms a solution with a high ph.
- reactivity increases down group
Trend in the halogens (group 17)
- very reactive group of non-metals
- they have characteristic colors. Cl= dense pale green, smelly and poisonous. Br= red/brown liquid, smelly and poisonous. Iodine= grey solid with purpule vapor, smelly and poisonous
- physical state changes down the group, with chlorine being a gas to bromine a liquid. This is related to the increasing mole mass which results in stronger intermolecular forces between molecules.
- halogens are diatomic
- melting/boiling points increase down the group, due to increase IMF
- reactivity decreases down group
- react with group 1 to form salts.
What is a displacement reaction?
A reaction in which the more reactive halogen displaces ions of the less reactive halogen from solution. They are accompanied by color changes
