periodic table.pptx
TRANSCRIPT
![Page 1: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/1.jpg)
Periodic table :periodicity
![Page 2: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/2.jpg)
The Modern Periodic Table
• The periodic table is an arrangement of the chemical elements, organized on the basis of their atomic numbers, electron configurations and recurring chemical properties.
• There are 92 naturally occurring elements. • The Modern Periodic Table is made up of 18
groups and 7 period.
![Page 3: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/3.jpg)
The modern periodic table
![Page 4: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/4.jpg)
The physical properties of
elements of period 2 and 3
![Page 5: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/5.jpg)
Atomic radius (atomic radii)
• The size of an atom cannot be measured exactly, However we can measure the size of atom in terms of its atomic radius.
• The atomic radius is half the distance between the nuclei of two closest and identical atoms
![Page 6: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/6.jpg)
![Page 7: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/7.jpg)
• We can classify atomic radius into threea) Covalent radiusb) Metallic radiusc) Van der Waals radius
![Page 8: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/8.jpg)
Covalent radius• Definition : half the length from one nuclei to
another atom’s nuclei bonded covalently.
![Page 9: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/9.jpg)
Metallic radius
• Definition : half of the distance between the two adjacent metal atoms in the metallic lattice.
![Page 10: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/10.jpg)
![Page 11: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/11.jpg)
NOTE #For metallic elements it may refer to covalent radius or metallic radius.
>Metallic used for non metal>Covalent for metals
![Page 12: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/12.jpg)
• For metallic elements which consists of covalent radius or metallic radius they have generally smaller atomic radius .
• Covalent bonds and lattice are very strong bonds that pull the shells closer together, causing the radius to decrease.
![Page 13: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/13.jpg)
Van der Waals radius• Definition : half the distance between two
neighbouring atoms which are not chemically bonded in solid state.
• Appear like touching, less attractive force.
![Page 14: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/14.jpg)
![Page 15: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/15.jpg)
Factors affecting Atomic radius
• There are factors that affect the atomic radius (size of atom ) which causes atoms having different sized across the group and down the period.
a) Screening effectb) Nuclear charge
![Page 16: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/16.jpg)
Screening effect
• The decrease in attraction between an electron and the nucleus of an atom with more then one electron.
• Its caused by mutual repulsion between electrons in the inner shell with those at the outer shell.
• This repulsion (screening ) causes the size of atom to increase.
![Page 17: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/17.jpg)
+
Mutual repulsion
![Page 18: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/18.jpg)
• The increase number of electrons the higher the repulsion force.
• The inner electrons shield the outer electrons from the nucleus pull
• greater the screening effect, easier the removal of electron.
![Page 19: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/19.jpg)
• However if the inner electronic shells have electrons , the attraction forces between nucleus and outermost electron will not be strong.
• The outermost electron is shield from the nucleus by the inner electrons.
![Page 20: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/20.jpg)
11+
![Page 21: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/21.jpg)
• Eg : magnesium atom has larger atomic radius then berylium. Although Mg has a larger nuclear charge it has more occupied shell then berylium.
• The more the number of electrons, less the attraction force , smaller the size
(refer periodic table )• The lower in the period the more screening
effect the larger the atomic radius.
![Page 22: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/22.jpg)
![Page 23: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/23.jpg)
Nuclear charge• The nucleus charge is the total charge of all the
protons in the nucleus. • It has the same value as the atomic number.• The nuclear charge increases as you go across the
periodic table. • The nuclear charge pulls the electrons closer to the
nucleus and causes the atomic size to decrease.• The stronger the nucleus charge (atomic number ) the smaller the atomic radius
![Page 24: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/24.jpg)
• Eg : across period 2 carbon has atomic number of 6 and nitrogen 7. this means nitrogen has higher nuclear charge hence smaller atomic radius.
• The effective nuclear charge : the difference between the screening constant and the actual nuclear charge.
• The higher the effective nuclear charge the smaller the atomic radius
![Page 25: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/25.jpg)
![Page 26: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/26.jpg)
Comparing =)
Screening effect• Down the group the size
increases• The number of shells with
occupied electrons increase• The attraction force
decreases• The size increases
Nuclear charge• Across the period the size
decreases• The proton number
increases across the period• The attraction force
increases• The size decreases
![Page 27: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/27.jpg)
![Page 28: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/28.jpg)
Atomic radii (radius ) across period 2 and 3
Across the period 2 and 3 (from left to right ) there is an decreases in atomic size.This is due to the increase in nuclear charge across these periodsHence increasing its electrostatic pull between electrons and nucleus, resulting in decrease in atomic size.
![Page 29: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/29.jpg)
• The screening effect will remain almost unchanged as the electrons across the same period are added to the same quantum shell which are 2s (period 2 ) and 2p( period 3).
• This will cause the effective nuclear charge to increase
• Due to the outermost electrons being pulled closer to the nucleus hence, decreasing its atomic radius.
![Page 30: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/30.jpg)
Atomic radii down a group.
• Going down a group the atomic radius will increase.
• This is because the increase in proton number that results in increase number of shells. This causes the attraction forces between nucleus and outermost electrons to decrease.
• Increase in screening effect ( repulsion )• This causes the atomic size to increase.
![Page 31: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/31.jpg)
![Page 32: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/32.jpg)
![Page 33: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/33.jpg)
Ionic radii
• The radius of a atom’s ion (cation or anion)• Ionic radii of the cation and anion gives the
distance between the ions in a crystal lattice. • The higher the nuclear charge, the higher the
forces of attraction and hence larger ionic radius
![Page 34: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/34.jpg)
![Page 35: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/35.jpg)
Ionic radii across period 2 and 3(The size of cations and anions decrease)
Cations decrease with increasing proton numbera) The increase in proton number increases the
attraction force between nucleus and electron hence smaller atomic size to decrease.
b) In period 2,all cations have valence of 2 but as the proton number increases as for Li+ (3) and (5) for B3+ so does the strength.
c) The attraction between B3+ is stronger then Li+
![Page 36: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/36.jpg)
The anion size decreases with increasing proton number
a) The proton number increasing causes the attraction between the nucleus and electrons to increase hence the atomic size decrease.b) All anions in period 2 has 8 electrons but the nuclear charge increases across the period causes the attraction to increase• Eg : N3+ (7) and F-(9)
![Page 37: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/37.jpg)
The ionic radius down the group
• Going down the group the size of ions increases
• The screening effect increases with the addition of extra shell.
• The more the number of shells the less the attraction force ( increase in repulsion ) hence the atomic size increases
![Page 38: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/38.jpg)
![Page 39: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/39.jpg)
• Neutral atoms or ions of same number of electrons are said to be isoelectronic.
• Positive ions (cations) are smaller than its neutral atom.
• Na+ ion is smaller than Na atom. This is because Na+ is more stable by donating one electron hence the attraction force is stronger
• Negative ions (anions) are larger than its neutral atoms .
• Cl- is larger than Cl atom as Cl- is more stable by accepting one electron making its attraction force between nucleus weaker
![Page 40: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/40.jpg)
Melting/boiling point and the enthalpy of vaporisation.
• Melting point is the temperature when a solid changes into liquid
• Boiling point is the temperature when the vapor pressure of the liquid is equal of that atmospheric pressure.
• Enthalpy of vaporisation is the heat energy required to covert 1mole of liquid into vapour at its boiling point.Also it measures the strength of the intermolecular forces between particles in its liquid state.
![Page 41: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/41.jpg)
• The melting points depend on two factors:a) The bonds involved (ionic, covalent , metallic
or Van der Waals )b) The structure/particle arrangements ( giant
covalent , simple molecular or metallic s)
![Page 42: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/42.jpg)
CHEMICAL BONDING• For elements with strong bonding like ionic
solids or giant covalent molecules their melting and boiling points are very high
• Because their molecules are held by strong attraction forces that makes it harder to break
• More energy is needed to break the bonds hence making the melting and boiling point higher
![Page 43: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/43.jpg)
MOLECULAR STRUCTURE
• A covalently bonded molecule between atoms in giant crystal lattice.
• Metallic bonds in metal lattice, the positively charges metal ions are attracted to the could of electron. As the valence electron increase the attraction force increases
• Simple molecular structure are for non metallic elements that form simple structure.
•
![Page 44: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/44.jpg)
![Page 45: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/45.jpg)
Based on the graph, • sodium, magnesium and aluminium are giant
molecules with metallic bonds• Silicon has giant covalent molecule.• Phosphorus , sulphur , chlorine and argon are
simple covalent molecule.
![Page 46: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/46.jpg)
• Hence the boiling point increases gradually for Na, Mg ,Al and Si as they have a strong covalent/metallic bond (intermolecular forces )that requires more energy to break the bonds causing a steep increase in boiling/melting point
• As for the simple molecules P,S,Cl they require less energy hence the boiling/melting point experience a decrease.
• The strength of metallic/covalent bond increases with the increase in valence electron.
![Page 47: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/47.jpg)
• For the enthalpy of vaporisation it is directly proportional to the melting/boiling.
• As the boiling/melting increases the enthaply of vaporisation also increases as more heat is needed to convert 1mole of liquid into vapour.
![Page 48: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/48.jpg)
Electrical Conductivity
![Page 49: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/49.jpg)
• Metals are good conductors of electricity .This is due to the presence of mobile electrons
• Non-metals do not conduct electricity.• As the number of delocalised electrons
increase the electric conductivity increases.
![Page 50: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/50.jpg)
![Page 51: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/51.jpg)
• Li and Be are metals that conduct good electricity in solid and molten state due to their delocalised electrons that move freely across the metal.
![Page 52: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/52.jpg)
Electronegativity
![Page 53: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/53.jpg)
• Is the ability of an atom in a covalent bond to attract shared electrons to itself
• Tendency to attract electrons.• The greater its electronegativity the greater its
tendency to attract electron.• Is effected by the >atomic number(nuclear charge) or>distance of valence electron from nucleus(atomic radius)
![Page 54: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/54.jpg)
The electronegativity increases across the period
• As the atomic radius decreases across the period its electronegativity increases across the period
• Its attraction forces between nucleus and electron will increase hence the tendency to attract electrons are higher
![Page 55: Periodic table.pptx](https://reader036.vdocuments.us/reader036/viewer/2022062519/55cf92cb550346f57b999536/html5/thumbnails/55.jpg)
Electronegativity decreases down the group.
• As going down the group the nuclear charge and screening effect increases but, the screening effect has a bigger increase than the nuclear charge
• Hence the atomic radius increases• Resulting in an decrease in effective nuclear
charge causing the attraction between nucleus and electrons to decrease.
• Causing the tendency to attract electrons to decrease.