6. Atomic Size (Atomic Radius)
Oxygen lies to the right of nitrogen in the periodic table.
So the size of oxygen atom is:
a. Bigger than nitrogen atom
b. Smaller than nitrogen atom
c. Equal to nitrogen atom
7. Atomic Size (Atomic Radius)- ACROSS A
PERIOD
Atomic number increases
Number of Protons increases
Nuclear charge increases
Nuclear force of attraction on outermost shell
increases
ATOMIC SIZE DECREASES
8. Across a
Period Number of Protons
3 4 5 6 7 8 9
2,1 2,2 2,3 2,4 2,5 2,6 2,7
Li Be B C N O F
Atomic Size :
DECREASING
9. Atomic Size (Atomic Radius)
Oxygen lies to the right of nitrogen in the periodic table.
So the size of oxygen atom is:
a. Bigger than nitrogen atom
b. Smaller than nitrogen atom
c. Equal to nitrogen atom
10. Atomic Size (Atomic Radius)- DOWN A
GROUP
Atomic number increases
Nuclear charge increases
Number of shells increase successively
(overweigh the increased nuclear charge)
Distance between the outermost shell and
nucleus increases
ATOMIC SIZE INCREASES
11. Down a group
atomic size:
H (1)
Li (2,1)
Na (2,8,1)
K (2, 8, 8, 1)
INCREASING
13. Atomic
Radius
Try this!
Referring to a periodic table, arrange the following
atoms in order of increasing atomic radius: P, Si,
N.
Answer:
N < P < Si
15. Ionization Energy
•the minimum energy (in kJ/mol) required to
remove an electron from a gaseous atom and
convert it into a positively gaseous ion.
M + I.E. M+ + e-
16. Across a
Period Number of Protons
3 4 5 6 7 8 9
2,1 2,2 2,3 2,4 2,5 2,6 2,7
Li Be B C N O F
IONIZATION INCREASES
17. Ionization Energy ACROSS A PERIOD
Atomic number increases
Nuclear charge increases
Atomic size decreases
Electrons are held tightly, hence greater energy
is required to remove the electrons
IONIZATION ENERGY INCREASES
21. Ionization Energy - DOWN A GROUP
Nuclear charge increases
Number of shells increase successively
(overweigh the increased nuclear charge)
Atomic size increases
Electrons are held loosely, hence lesser energy
is required to remove the electrons
Ionization Energy decreases
22. IRREGULARITIES ON IONIZATION ENERGY
The first exception occurs between Group 2A and 3A
elements in the same period (for example, between Be and B
and between Mg and Al).
The second irregularity occurs between Groups 5A and 6A
(for example, between N and O and between P and S).
24. Ionization Energy
The increase in first ionization energy from left to right
across a period and from bottom to top in a group for
representative elements.
27. Across a
Period Number of Protons
3 4 5 6 7 8 9
2,1 2,2 2,3 2,4 2,5 2,6 2,7
Li Be B C N O F
Electronegativity:
INCREASING
28. Electronegativity ACROSS A PERIOD
Atomic number increases
Nuclear charge increases
Atomic size decreases
Tendency to attract the shared pair of electrons
increases
ELECTRONEGATIVITY INCREASES
31. Electronegativity- DOWN A GROUP
Nuclear charge increases
Number of shells increase successively
(overweigh the increased nuclear charge)
Atomic size increases
Tendency to attract the shared pair of electrons
decreases
ELECTRONEGATIVITY DECREASES
33. Electron Affinity (E.A)
The amount of energy released while
converting neutral gaseous atom into a
negatively charged gaseous ion by the
addition of electron
X + e- X- + EA
34. Across a
Period Number of Protons
3 4 5 6 7 8 9
2,1 2,2 2,3 2,4 2,5 2,6 2,7
Li Be B C N O F
Electron Affinity:
INCREASING
35. Electron Affinity ACROSS A PERIOD
Atomic number increases
Nuclear charge increases
Atomic size decreases
Effective force of attraction on valence electrons
increases, hence greater energy is released
ELECTRON AFFINITY INCREASES
38. Electron Affinity- DOWN A GROUP
Nuclear charge increases
Number of shells increase successively
(overweigh the increased nuclear charge)
Atomic size increases
Effective force of attraction on valence electrons
decreases, hence lesser energy is released
ELECTRON AFFINITY DECREASES
39. IRREGULARITIES ON ELECTRON AFFINITY
The electron affinity of a Group 2A element is
lower than that for the corresponding Group 1A
element, and the electron affinity of a Group 5A
element is lower that that for the corresponding
Group 4A element.
Editor's Notes
One important property of the element is size.
Calculatead as the Distance between the center nucleus to the outermost shell. The distance is the atomic radius.
Sometimes atoms don’t exist independently. They form molecules, in that case, we take the internuclear distance. That is the distance between the nucleus of the 2 atoms. And this distance is then half . Then gives the particular radius of the atom.
A number of physical properties, including density, melting point, and boiling point, are related to the sizes of atoms, but atomic size is difficult to define. the electron density in an atom extends far beyond the nucleus, but we normally think of atomic size as the volume containing about 90 percent of the total electron density around the nucleus.
we define the size of an atom in terms of its atomic radius, which is one-half the distance between the two nuclei in two adjacent metal atoms.
INCREASING NUMBER OF PROTONS PULLS THE ELECTRON FROM THE OUTERMOST SHELL CLOSER TO ITSELF. THEREFORE, DECREASING THE SIZE OF THE ATOM.
INCREASING NUMBER OF PROTONS PULLS THE ELECTRON FROM THE OUTERMOST SHELL CLOSER TO ITSELF. THEREFORE, DECREASING THE SIZE OF THE ATOM.
Atomic radii (in picometers) of representative elements according to their positions in the periodic table. Note that there is no general agreement on the size of atomic radii. We focus only on the trends in atomic radii, not on their precise values.
Consider the second-period elements. Because the effective nuclear charge increases from left to right, the added valence electron at each step is more strongly attracted by the nucleus than the one before. Therefore, we expect and indeed fi nd the atomic radius decreases from Li to Ne. Within a group, we fi nd that atomic radius increases with atomic number.
GOING DOWN A GROUP: INCREASING ENERGY LEVEL : INCREASING RADIUS
ACROSS A PERIOD LEFT TO RIGHT: INCREASING ELECTRONS, THE MORE IT GETS ATTRACTED, THE CLOSER THE ELECTRONS TO NUCLEUS: SMALLER RADIUS
Atomic radii (in picometers) of representative elements according to their positions in the periodic table. Note that there is no general agreement on the size of atomic radii. We focus only on the trends in atomic radii, not on their precise values.
Consider the second-period elements. Because the effective nuclear charge increases from left to right, the added valence electron at each step is more strongly attracted by the nucleus than the one before. Therefore, we expect and indeed fi nd the atomic radius decreases from Li to Ne. Within a group, we fi nd that atomic radius increases with atomic number. For the alkali metals in Group 1A, the valence electron resides in the ns orbital. Because orbital size increases with the increasing principal quantum number n, the size of the atomic radius increases from Li to Cs.
Atomic radii (in picometers) of representative elements according to their positions in the periodic table. Note that there is no general agreement on the size of atomic radii. We focus only on the trends in atomic radii, not on their precise values.
Consider the second-period elements. Because the effective nuclear charge increases from left to right, the added valence electron at each step is more strongly attracted by the nucleus than the one before. Therefore, we expect and indeed fi nd the atomic radius decreases from Li to Ne. Within a group, we fi nd that atomic radius increases with atomic number. For the alkali metals in Group 1A, the valence electron resides in the ns orbital. Because orbital size increases with the increasing principal quantum number n, the size of the atomic radius increases from Li to Cs.
In other words, ionization energy is the amount of energy in kilojoules needed to strip 1 mole of electrons from 1 mole of gaseous atoms. / CATIONIZATION
The higher the ionization energy, the more difficult it is to remove the electron.
The smaller the size, the more difficult it is for the electrons to be removed: higher ionization energy
L-R= INCREASING ie
T-B= DECREASING IE ( the farther energy level, causes electron repulsion, weaker bond, so easier to be removed: low IE REQUIRED
INCREASING NUMBER OF PROTONS PULLS THE ELECTRON FROM THE OUTERMOST SHELL CLOSER TO ITSELF. THEREFORE, DECREASING THE SIZE OF THE ATOM.
IE1 AMOUNT OF ENERGY REQUIRED WHEN REMOVING THE FIRST ELECTRON, IE2 IS THE SECOND
As more electron is removed, the smaller the size, the more difficult it is to remove the electrons, thus it requires high IE
INCREASING NUMBER OF PROTONS PULLS THE ELECTRON FROM THE OUTERMOST SHELL CLOSER TO ITSELF. THEREFORE, DECREASING THE SIZE OF THE ATOM.
The Group 3A elements have lower first ionization energies than 2A elements because they all have a single electron in the outermost p subshell (ns2np1), which is well shielded by the inner electrons and the ns2 electrons. Therefore, less energy is needed to remove a single p electron than to remove a paired s electron from the same principal energy level.
In the Group 5A elements (ns2np3) the p electrons are in three separate orbitals according to Hund’s rule. In Group 6A (ns2np4) the additional electron must be paired with one of the three p electrons. The proximity of two electrons in the same orbital results in greater electrostatic repulsion, which makes it easier to ionize an atom of the Group 6A element, even though the nuclear charge has increased by one unit. Thus, the ionization energies for Group 6A elements are lower than those for Group 5A elements in the same period.
Atomic radii (in picometers) of representative elements according to their positions in the periodic table. Note that there is no general agreement on the size of atomic radii. We focus only on the trends in atomic radii, not on their precise values.
Consider the second-period elements. Because the effective nuclear charge increases from left to right, the added valence electron at each step is more strongly attracted by the nucleus than the one before. Therefore, we expect and indeed fi nd the atomic radius decreases from Li to Ne. Within a group, we fi nd that atomic radius increases with atomic number. For the alkali metals in Group 1A, the valence electron resides in the ns orbital. Because orbital size increases with the increasing principal quantum number n, the size of the atomic radius increases from Li to Cs.
Atomic radii (in picometers) of representative elements according to their positions in the periodic table. Note that there is no general agreement on the size of atomic radii. We focus only on the trends in atomic radii, not on their precise values.
Consider the second-period elements. Because the effective nuclear charge increases from left to right, the added valence electron at each step is more strongly attracted by the nucleus than the one before. Therefore, we expect and indeed fi nd the atomic radius decreases from Li to Ne. Within a group, we fi nd that atomic radius increases with atomic number.
GOING DOWN A GROUP: INCREASING ENERGY LEVEL : INCREASING RADIUS
ACROSS A PERIOD LEFT TO RIGHT: INCREASING ELECTRONS, THE MORE IT GETS ATTRACTED, THE CLOSER THE ELECTRONS TO NUCLEUS: SMALLER RADIUS
Which can pull the shared electrons?
Atomic radii (in picometers) of representative elements according to their positions in the periodic table. Note that there is no general agreement on the size of atomic radii. We focus only on the trends in atomic radii, not on their precise values.
Consider the second-period elements. Because the effective nuclear charge increases from left to right, the added valence electron at each step is more strongly attracted by the nucleus than the one before. Therefore, we expect and indeed fi nd the atomic radius decreases from Li to Ne. Within a group, we fi nd that atomic radius increases with atomic number.
GOING DOWN A GROUP: INCREASING ENERGY LEVEL : INCREASING RADIUS
ACROSS A PERIOD LEFT TO RIGHT: INCREASING ELECTRONS, THE MORE IT GETS ATTRACTED, THE CLOSER THE ELECTRONS TO NUCLEUS: SMALLER RADIUS
Smaller size: increased attraction of e- towards the nucleus: increasing ELECTRONEGATIVITY
Atomic radii (in picometers) of representative elements according to their positions in the periodic table. Note that there is no general agreement on the size of atomic radii. We focus only on the trends in atomic radii, not on their precise values.
Consider the second-period elements. Because the effective nuclear charge increases from left to right, the added valence electron at each step is more strongly attracted by the nucleus than the one before. Therefore, we expect and indeed fi nd the atomic radius decreases from Li to Ne. Within a group, we fi nd that atomic radius increases with atomic number. For the alkali metals in Group 1A, the valence electron resides in the ns orbital. Because orbital size increases with the increasing principal quantum number n, the size of the atomic radius increases from Li to Cs.
INCREASING NUMBER OF PROTONS PULLS THE ELECTRON FROM THE OUTERMOST SHELL CLOSER TO ITSELF. THEREFORE, DECREASING THE SIZE OF THE ATOM.
Which can pull the shared electrons?
Another property that greatly infl uences the chemical behavior of atoms is their ability to accept one or more electrons
This property is called electron affi nity, which is the negative of the energy change that occurs when an electron is accepted by an atom in the gaseous state to form an anion.
Gaseous – isolated atoms, not associated to other atoms
Greater Attraction of electron, greater energy released
The halogens (Group 7A) have the highest electron affi nity values.
Decreasing size: greater attraction
Outermost shell becomes closer : the attraction of e- to the nucleus increases
GREATER AFFINITY: GREATER ENERGY RELEASED
Atomic radii (in picometers) of representative elements according to their positions in the periodic table. Note that there is no general agreement on the size of atomic radii. We focus only on the trends in atomic radii, not on their precise values.
Consider the second-period elements. Because the effective nuclear charge increases from left to right, the added valence electron at each step is more strongly attracted by the nucleus than the one before. Therefore, we expect and indeed fi nd the atomic radius decreases from Li to Ne. Within a group, we fi nd that atomic radius increases with atomic number. For the alkali metals in Group 1A, the valence electron resides in the ns orbital. Because orbital size increases with the increasing principal quantum number n, the size of the atomic radius increases from Li to Cs.
INCREASING NUMBER OF PROTONS PULLS THE ELECTRON FROM THE OUTERMOST SHELL CLOSER TO ITSELF. THEREFORE, DECREASING THE SIZE OF THE ATOM.
These exceptions are due to the valence electron confi gurations of the elements involved. An electron added to a Group 2A element must end up in a higher-energy np orbital, where it is effectively shielded by the ns2 electrons and therefore experiences a weaker attraction to the nucleus. Therefore, it has a lower electron affi nity than the corresponding Group 1A element. Likewise, it is harder to add an electron to a Group 5A element (ns2np3) than to the corresponding Group 4A element (ns2np2) because the electron added to the Group 5A element must be placed in a np orbital that already contains an electron and will therefore experience a greater electrostatic repulsion. Finally, in spite of the fact that noble gases have high effective nuclear charge, they have extremely low electron affi nities (zero or negative values). The reason is that an electron added to an atom with an ns2np6 confi guration has to enter an (n 1 1)s orbital, where it is well shielded by the core electrons and will only be very weakly attracted by the nucleus. This analysis also explains why species with complete valence shells tend to be chemically stable.
