Ch7 Trends In The Periodic Table Flashcards
Heisenbergs uncertainty principle
Not possible to measure the distance between the nucleus and the outermost electrons of an atom
Atomic radius
Of an atom is defined as half the distance between the nucleus and of two atoms of the Sam element that are joined together by a single covalent bond
Bond length
How is atomic radius measured
X-ray diffraction and electron diffraction
What is the covalent radius of a hydrogen molecule with distance between two nucleus = 0.074nm
0.037nm
Why do noble gasses not have value for covalent radius
They don’t form covalent bonds with one another
Atomic radius trend
Decrease across periods/rows
Increase down groups
What is the size of an atom governed by
The electrostatic attraction between the positively charged protons and negatively charged electrons
If electrostatic attraction between protons and electrons is large
the positive protons will pull the outer electrons closer to the nucleus giving a smaller atomic radius
If electrostatic attraction is small
The electrons will be further from the nucleus giving a larger atomic radius
Reasons: atomic radius increase down the groups
- New energy level/ shells - additional electrons go into new energy levels as you go down. Outer electrons become further away from nucleus
- Screening effect of inner electrons - electrons in inner energy levels shield/ screen outer electrons from positive charger of the nucleus giving a
Reason: atomic radius decrease across a period
- Increase in Effective Nuclear Charge - no. Of protons in nucleus increase from left to right in periodic table. Greater attractive force on outer electrons, draws them closer to nucleus
- no increase in screening effect - extra electrons go into same outer shell as you go across
First ionisation energy
Of an atom is the minimum energy required to completely remove the most loosely bound electrons of from a neutral gaseous atom in its ground state
Na(g) ->Na+(g) + e-
First ionisation energy decrease down a group
- Increasing atomic radius
- Screening effect of inner electrons
First ionisation increases across a period
- Increasing effective nuclear charge
- Decreasing atomic radius increase
Exceptions to general trend across a period
-Higher than expected ionisation energies
- extra stability filled or half- filled sub levels
(Look to work copy)
Second ionisation energy
Energy required to remove an electron from an ion with one positive charge in the gaseous state
X+(g) -> X2+(g) + e-
Elevtronegeativity
The relative attraction that an atom in a molecule has for the shared pair of electrons in a covalent bond
General points ionistaion energies
- steady increase as electrons removed, ion becomes more positive, greater attraction on remaining electrons
-large increase when electron removed from new energy level/ shells
-substantial increase when e removed from new sublevel
Electro negativity trends
Decreases down a group
Increases across a period
Electronegativity decrease down the groups
- Increasing atomic radius
-outer electrons are becoming further away from attractive force of nucleus
-smaller attraction between nucleus and the shared pair of electrons … electroneg decreases - Screening effect of inner electrons
-even though nuclear charge increases down groups, is cancelled out by the screening effect of intervening shells of electrons
-outer electrons are shielded from attractive force of pos charged nucleus
-attraction force of nucleus/ electronegativity decreases down groups
Electronegativity increase across the periods
- Increasing effective nuclear charge
-moving left to right, number of protons in the nucleus increases, therefore attraction between nucleus and outer electrons is also increasing - electrons involved in binding are more strongly attracted to the nucleus - Decreasing atomic radius
- within any row, the atomic radius decreases from left to right. Therefore outer electrons are closer to nucleus.
Thus there is greater attraction between nucleus and the outer electrons
Group 1
Alkali metals
Increasing reactivity down the group
Group 1 points
-very reactive, low first ionisation values
-none occur free in nature, stored under oil or in glass container with air removed (Rb, Cs)
-low melting points
-readily lose single outer electron to form ionic compounds
Group 1 chemical reactions with oxygen
Alkali metals react with oxygen to form oxides
Eg
Potassium + oxygen -> potassium oxide
2k + 1/2 O -> K2O
Chemical reactions with water - group 1
Alkali metals react with water to form hydroxide of metal and hydrogen gas is given off
Eg
Sodium + water -> sodium hydroxide + hydrogen
Na + H2O -> NaOH + 1/2H2
Group 1 chemical reactions with acid
Too dangerous, explosive
Sodium + hydrochloric acid -> sodium + hydrogen chloride
Na + HCL -> NaCl + 1/2H
Group 7
Halogens
Increasing reactivity up the group
Properties of halogens group 7
-most electronegative elements in table
-fluorine most electronegative element, Electronegativity decreases down group
-don’t exist free in nature
-oxidising agents - removes electrons from other substances easily
Trend in chemical reactivity of halogens 1 - chlorine + bromide
Chlorine gas bubbles through solution of bromide ions, chlorine takes electrons from bromide
Chlorine displaces bromide from solution
Chlorine + bromide —> chloride ion + bromine
Cl2 + 2Br- —> 2Cl- + Br2
Trend in chemical reactivity of halogens 2 - bromine + iodine
Bromine displaces iodine from solution
Bromine + iodine ion —> bromide ions + iodine
Br2 + 2I- —> 2Br- + I2
The more reactive halogens displaces the less reactive halogen from the solution of its ions
Group 2
Alkaline earth metals
Group 2 alkaline earth metals react with water and reason
Be does not
Magnesium reacts very slowly
Calcium undergoes steady reaction
Reason: ionisation energies decrease down the group
Strontium and barium react more vigorously with water to form
Noble inert gases properties
Form practically no compounds, unreactive.
-steady increase in boiling points down the group he - radon
(Increasing atomic radius… more van der waals forces)
Boiling points of halogens
Increase down the group
-due to stronger van der waals forces as the atomic radius increases down the group
Why do half full and full outer sub levels have higher ionisation energies
Extra stability, more difficult to remove electron and causes requires sudden increase in ionisation energy
What happens. As you take away electrons
Atom becomes more positive and is holding into elect Ron’s more tightly
-full outer sublevel more stable
