Periodicity Flashcards
How are elements arranged on the periodic table
In order of increasing atomic number
Horizontal rows are called
Periods
As you move along a period what happens to each element
Increases by a proton and an electron, hence an increase in atomic number
Vertical columns are called what
GroupsW
What is special about elements in the same group
Have similar chemical properties
How is the periodic table divided in terms of blocks
According to electronic sub-level being filled e.g. S, P, D, F blocks
What block are groups 1 and 2
S block, electrons are being added to the S sub-shell
What blocks are groups 3 to 7 and 0
P block, electrons are being added to the P subshell
What elements on the Periodic table consist of D block
Transition metals
What element on the periodic table are the F blocks
Lanthanides and the Actinides
What determines the position and chemical properties of elements
Their electronic configuration
How can you predict properties of elements
By seeing where they are positioned on the periodic table
Define Periodicity
A repeating trend in the properties of elements across each period of the periodic table
What properties display periodicity?
-ionization energy
-atomic radius
-melting points
-boiling points
What is atomic radius?
Distance from the center of the nucleus to the outermost electrons (valence electrons)
How can you calculate atomic radii
Covalent radii
Metallic Radii
How do you calculate atomic radii using Covalent radii?
-calculated by measuring distance between two covalently bonded atoms and dividing by two to get individual radii (atoms have to be the same element otherwise their diameters will not be equal
How do you calculate atomic radii using Metallic radii?
Metallic radii, calculated by measuring diameter of a cluster of metal atoms and then dividing by the number of atoms
-only works for atoms of the same type
As you move across a period what happens to the atomic radii?
Decreases
Why does atomic radii decrease across a period?
As you go across a period:
- a proton is added to nucleus= increases nuclear charge= greater pull generated from nucleus to all e-
- Shielding remains constant
- meaning all e- are drawn closer to nucleus
-T/F less distance between valence e- and center of nucleus (atomic radii)
Define First Ionization energy.
The energy required to remove one electron from each atom in one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions.
As you move across period 3 what happens to the Ionization energy?
Increases
Why does Ionization energy increase across period 3?
-atomic radii decreases
-shielding remains constant
-no. of protons is increasing so positive charge on nucleus increases
-causes electrons to be held more tightly so more energy is needed to release them
As you move across period 3 is there a steady increase in IE?
-not a steady rise because of different sublevel
- Na+ Mg= increasing
-DIP at Al
- Al+ Si+ P= increasing
- DIP at S
-S+ Cl +Ar= increasing
Why is there not a steady increase in IE as you move across period 3? (5 marks)
- because of different sublevels
- First rise= for the s orbital
-A dip for Al= electrons are being removed from p orbital where electrons are at higher energies so easier to remove
-Next rise= for first 3 electrons of p orbitals where unpaired electrons are being removed
- 2nd Dip= when first paired electron is removed in S, there is more electrostatic repulsion between paired electrons so they are easier to remove
-Trend then continues to increase as expected due to increase in Nuclear charge
Define Electronegativity.
A measure of the attraction of a bonded atom for the pair of electrons in a covalent bond
The more electronegative an atom is…
…The more strongly it pulls electrons towards itself
What can electronegativity be used to predict?
The polarity of a covalent bond
First ionization energy is…
energy needed to remove 1 electron from each atom in 1 mole of gaseous atoms
Look at periods 2 and 3 only. What happens to the size of the atoms as we move across a period? Why?
Atoms get smaller across a period: atomic radii decreases
-e- shielding remains constant
-number of outer shells stays the same
-the number of protons increases across a period so the positive nuclear charge increases meaning nuclear attraction increases
-as electrons are being held more tightly
-decreasing atomic radii- distance between nucleus and valence electrons
How will the atomic radius decreasing have an affect on the First Ionization energy?
First IE increases-
-atomic radius decreases
-electrons are being held more closely to nucleus as nuclear charge increases pulling the electrons
-electron shielding stays the same
Overall there is an increase to First IE as the tendency to loose electrons on an atom decreases, because more electrons are on the same outer shell and have stronger attraction to nucleus
What happens to atomic radius, nuclear charge and electron shielding as we move across period 2 and 3?
Atomic radius= decreases
Nuclear charge= Increases
Electron shielding= remains constant as electrons are on the same outer shell
Overall increase in nuclear attraction as electrons are being held more closely to the nucleus, needing more energy to remove them
What happens to electron shielding as we move across a period?
No. of shells present stays the same= electron shielding remains constant
Electron shielding remains constant across periods 2 and 3. What affect will this have on strength of attraction between nucleus and valence electrons?
- electron shielding remains constant but nuclear charge increases:
-as we move across a period, more protons are being added onto the nucleus of each atom
-as we move across a period there are also increasing electrons in each atom
-Electrons stay in the same outer shell- these are attracted to the positive nucleus which pulls the valence electrons decreasing atomic radius
-over all increase in nuclear attraction
What happens to the nuclear charge as we move across a period? How does this effect the first ionization energy?
Number of protons increases as we move across a period, therefore there is a stronger nuclear attraction between electrons and nucleus
-This increases the first ionization energy
Describe the trend across a period
Number of electrons increases
Number of protons increases
Nuclear charge increases
Atomic radius decreases
Electron Shielding remains constant
Nuclear attraction increases
Increases First Ionization energy
What happens to the size of any atoms as we move down a group?
As we move down a group the atoms become larger
As we move down a group the atoms become larger. What affect will this have on the first ionization energy?
Atomic radius increases
-Distance between valence electrons and nucleus increases, so we need less energy to remove it as nuclear attraction is lost
What happens to electron shielding as move down a group?
Number of electron shells increases so amount of electron shielding increases
Electron shielding increases down a group, how does this affect the strength of the attraction between valence electrons and nucleus?
-no of shells present increases
-electron shielding increases
- Nuclear attraction is weakened between valence electrons and positive nucleus due to more shielding
What happens to the nuclear charge as we move down a group?
Nuclear charge increases, down a group however its effect is outweighed by increased atomic radius and ( to a lesser extent) electron shielding
What has the bigger effect in terms of trends down a group?
-Increasing atomic radii
-increasing electron shielding (to a lesser extent)
What is the trend as we move down any group?
-atomic radius increases
-electron shielding increases
-Nuclear charge increases= effect is outweighed by atomic radius
-Nuclear attraction on outer electron decreases
First IE decreases
Why are ionization values always positive?
Because energy must be supplied (an endothermic energy change) to seperate electrons from atoms
What is successive ionization energy/ second ionization energy?
