14. what is electronegativity?
15. how does the electronegativity change going across a period? why?
16. how does the electronegativity change going down a group? why?
17. for each of the following pairs, circle the element with the greatest electronegativity.
a. y or rb
b. li or be
b. b or c
c. mg or sr
d. na or mg
f. k or na
g. be or b
h. n or o
i. ba or ra
j. ne or f
18. what is ionization energy?
19. how does the ionization energy change going across a period? why?
20. how does the ionization energy change going down a group? why?
21. for each of the following pairs, circle the element with the larger 1st ionization energy.
a. y or rb
b. li or be
c. b or c
d. mg or sr
e. na or mg
f. k or na
g. be or b
h. n or o
i. ba or ra
j. ne or f
23. which metal in the pairs below is more reactive? why?
_______ ba or ca
_______ li or na
_______ co or mt
24. which nonmetal in the pairs below is more reactive? why?
_______ f or cl
_______ se or s
_______ i or br

Question

14. what is electronegativity?
15. how does the electronegativity change going across a period? why?
16. how does the electronegativity change going down a group? why?
17. for each of the following pairs, circle the element with the greatest electronegativity.
    a. y or rb
    b. li or be
    b. b or c
    c. mg or sr
    d. na or mg
    f. k or na
    g. be or b
    h. n or o
    i. ba or ra
    j. ne or f
18. what is ionization energy?
19. how does the ionization energy change going across a period? why?
20. how does the ionization energy change going down a group? why?
21. for each of the following pairs, circle the element with the larger 1st ionization energy.
    a. y or rb
    b. li or be
    c. b or c
    d. mg or sr
    e. na or mg
    f. k or na
    g. be or b
    h. n or o
    i. ba or ra
    j. ne or f
23. which metal in the pairs below is more reactive? why?
    _______ ba or ca
    _______ li or na
    _______ co or mt
24. which nonmetal in the pairs below is more reactive? why?
    _______ f or cl
    _______ se or s
    _______ i or br

Accepted Answer

### Question 14: What is electronegativity?
# Brief Explanations:
Electronegativity is a chemical property that describes the tendency of an atom or a functional group to attract electrons (or electron density) towards itself in a chemical bond. It helps in predicting the nature of chemical bonds (ionic, covalent, polar covalent) and the distribution of electron density in molecules.
# Answer:
Electronegativity is the tendency of an atom to attract electrons in a chemical bond.

### Question 15: How does electronegativity change across a period? Why?
# Brief Explanations:
- **Trend**: Electronegativity generally increases as we move from left to right across a period (in the periodic table).
- **Reason**: As we move across a period, the number of protons in the nucleus (nuclear charge) increases while the number of electron shells remains the same. This increased nuclear charge exerts a stronger pull on the valence electrons, making the atom more electronegative (more capable of attracting electrons). Also, the atomic radius decreases across a period, so the valence electrons are closer to the nucleus, experiencing a stronger attractive force.
# Answer:
Electronegativity increases across a period. Reason: Increasing nuclear charge (more protons) and decreasing atomic radius pull valence electrons more strongly.

### Question 16: How does electronegativity change down a group? Why?
# Brief Explanations:
- **Trend**: Electronegativity generally decreases as we move down a group (in the periodic table).
- **Reason**: As we move down a group, the number of electron shells increases (atomic radius increases). The valence electrons are further away from the nucleus, and the shielding effect (due to inner electron shells) increases. This reduces the effective nuclear charge experienced by the valence electrons, so the atom's ability to attract electrons (electronegativity) decreases.
# Answer:
Electronegativity decreases down a group. Reason: Increasing atomic radius and shielding effect reduce effective nuclear charge on valence electrons.

### Question 17: For each pair, circle the element with the greatest electronegativity.
#### a. Y or Rb
- Y (Yttrium) and Rb (Rubidium) are in the same period (period 5). As electronegativity increases across a period, Y (which is to the right of Rb) has greater electronegativity.
#### b. Li or Be
- Li (Lithium) and Be (Beryllium) are in period 2. Electronegativity increases across a period, so Be (right of Li) has greater electronegativity.
#### c. B or C
- B (Boron) and C (Carbon) are in period 2. Electronegativity increases across a period, so C (right of B) has greater electronegativity.
#### d. Na or Mg
- Na (Sodium) and Mg (Magnesium) are in period 3. Electronegativity increases across a period, so Mg (right of Na) has greater electronegativity.
#### e. Mg or Sr
- Mg (Magnesium) and Sr (Strontium) are in group 2. Electronegativity decreases down a group, so Mg (above Sr) has greater electronegativity.
#### f. K or Na
- K (Potassium) and Na (Sodium) are in group 1. Electronegativity decreases down a group, so Na (above K) has greater electronegativity.
#### g. Be or B
- Be (Beryllium) and B (Boron) are in period 2. Electronegativity increases across a period, so B (right of Be) has greater electronegativity.
#### h. N or O
- N (Nitrogen) and O (Oxygen) are in period 2. Electronegativity increases across a period, so O (right of N) has greater electronegativity.
#### i. Ba or Ra
- Ba (Barium) and Ra (Radium) are in group 2. Electronegativity decreases down a group, so Ba (above Ra) has greater electronegativity.
#### j. Ne or F
- Ne (Neon) is a noble gas with a full valence shell, so it has very low electronegativity (it doesn't attract electrons as it's stable). F (Fluorine) is highly electronegative. So F has greater electronegativity.

# Answer:
a. $\boldsymbol{Y}$
b. $\boldsymbol{Be}$
c. $\boldsymbol{C}$
d. $\boldsymbol{Mg}$
e. $\boldsymbol{Mg}$
f. $\boldsymbol{Na}$
g. $\boldsymbol{B}$
h. $\boldsymbol{O}$
i. $\boldsymbol{Ba}$
j. $\boldsymbol{F}$

### Question 18: What is ionization energy?
# Brief Explanations:
Ionization energy (specifically first ionization energy) is the minimum amount of energy required to remove the most loosely bound electron (the valence electron) from an isolated gaseous atom to form a gaseous ion with a +1 charge (a cation), i.e., $X(g)
ightarrow X^+(g) + e^-$. It is a measure of how strongly an atom holds its electrons.
# Answer:
Ionization energy is the energy needed to remove a valence electron from a gaseous atom to form a +1 gaseous ion.

### Question 19: How does ionization energy change across a period? Why?
# Brief Explanations:
- **Trend**: Ionization energy generally increases as we move from left to right across a period.
- **Reason**: As we move across a period, the nuclear charge increases (more protons) while the atomic radius decreases (valence electrons are closer to the nucleus). The increased nuclear charge exerts a stronger pull on the valence electrons, so more energy is required to remove an electron. Also, the shielding effect remains relatively constant across a period (same number of inner shells), so the effective nuclear charge on valence electrons increases.
# Answer:
Ionization energy increases across a period. Reason: Increasing nuclear charge and decreasing atomic radius make it harder to remove electrons.

### Question 20: How does ionization energy change down a group? Why?
# Brief Explanations:
- **Trend**: Ionization energy generally decreases as we move down a group.
- **Reason**: As we move down a group, the atomic radius increases (more electron shells), and the shielding effect increases (inner shells shield valence electrons from the nucleus). The valence electrons are further from the nucleus, so the effective nuclear charge experienced by them is reduced. Thus, less energy is required to remove a valence electron.
# Answer:
Ionization energy decreases down a group. Reason: Increasing atomic radius and shielding effect reduce the hold on valence electrons.

### Question 21: For each pair, circle the element with the larger 1st ionization energy.
#### a. Y or Rb
- Y (Yttrium) and Rb (Rubidium) are in period 5. Ionization energy increases across a period, so Y (right of Rb) has larger 1st ionization energy.
#### b. Li or Be
- Li (Lithium) and Be (Beryllium) are in period 2. Ionization energy increases across a period, so Be (right of Li) has larger 1st ionization energy.
#### c. B or C
- B (Boron) and C (Carbon) are in period 2. Ionization energy increases across a period, so C (right of B) has larger 1st ionization energy.
#### d. Mg or Sr
- Mg (Magnesium) and Sr (Strontium) are in group 2. Ionization energy decreases down a group, so Mg (above Sr) has larger 1st ionization energy.
#### e. Na or Mg
- Na (Sodium) and Mg (Magnesium) are in period 3. Ionization energy increases across a period, so Mg (right of Na) has larger 1st ionization energy.
#### f. K or Na
- K (Potassium) and Na (Sodium) are in group 1. Ionization energy decreases down a group, so Na (above K) has larger 1st ionization energy.
#### g. Be or B
- Be (Beryllium) and B (Boron) are in period 2. Ionization energy generally increases across a period, so B (right of Be) has larger 1st ionization energy? Wait, no: Be has a filled 2s² subshell (stable configuration), while B has 2s²2p¹. Removing an electron from Be (a stable subshell) requires more energy than from B (which can lose the 2p¹ electron more easily). So Be has larger 1st ionization energy.
#### h. N or O
- N (Nitrogen) has a half - filled 2p³ subshell (stable configuration), O has 2p⁴. Removing an electron from N (stable half - filled subshell) is harder than from O. So N has larger 1st ionization energy? Wait, no: Wait, the general trend is increase across a period, but there are exceptions. N: 1s²2s²2p³ (half - filled p), O: 1s²2s²2p⁴. The electron - electron repulsion in O's 2p⁴ (paired electron) makes it easier to remove an electron from O than from N. So N has larger 1st ionization energy? Wait, actually, the first ionization energy of N is higher than O. So N has larger 1st ionization energy.
#### i. Ba or Ra
- Ba (Barium) and Ra (Radium) are in group 2. Ionization energy decreases down a group, so Ba (above Ra) has larger 1st ionization energy.
#### j. Ne or F
- Ne (Neon) is a noble gas with a full valence shell (very stable). F (Fluorine) has 7 valence electrons. Removing an electron from Ne (stable full shell) requires a huge amount of energy, much more than from F. So Ne has larger 1st ionization energy.

# Answer:
a. $\boldsymbol{Y}$
b. $\boldsymbol{Be}$
c. $\boldsymbol{C}$
d. $\boldsymbol{Mg}$
e. $\boldsymbol{Mg}$
f. $\boldsymbol{Na}$
g. $\boldsymbol{Be}$
h. $\boldsymbol{N}$
i. $\boldsymbol{Ba}$
j. $\boldsymbol{Ne}$

### Question 23: Which metal in the pairs is more reactive? Why?
#### Ba or Ca
- Ba and Ca are in group 2. Reactivity of metals in a group increases down the group (as ionization energy decreases, so they lose electrons more easily). Ba is below Ca in group 2, so Ba is more reactive.
#### Li or Na
- Li and Na are in group 1. Reactivity of metals in a group increases down the group (ionization energy decreases down the group, so they lose electrons more easily). Na is below Li in group 1, so Na is more reactive.
#### Co or Mt
- Co (Cobalt) is a transition metal, Mt (Meitnerium) is a very heavy element in the d - block (group 9, period 7). As we go down a group in transition metals, reactivity generally increases (due to lower ionization energy). Mt is below Co, so Mt is more reactive.

# Brief Explanations:
- For group 1 and 2 metals, reactivity increases down the group because ionization energy decreases (easier to lose electrons). For transition metals, reactivity generally increases down a group as well (lower ionization energy allows easier electron loss).
# Answer:
Ba (more reactive than Ca: group 2, reactivity increases down).
Na (more reactive than Li: group 1, reactivity increases down).
Mt (more reactive than Co: transition metal, reactivity increases down group).

### Question 24: Which nonmetal in the pairs is more reactive? Why?
#### F or Cl
- F and Cl are in group 17 (halogens). Reactivity of nonmetals (halogens) decreases down a group (as electronegativity decreases, so they gain electrons less easily). F is above Cl in group 17, so F is more reactive.
#### Se or S
- Se and S are in group 16. Reactivity of nonmetals in a group decreases down the group (electronegativity decreases down the group, so they gain electrons less easily). S is above Se in group 16, so S is more reactive.
#### I or Br
- I and Br are in group 17 (halogens). Reactivity of nonmetals (halogens) decreases down a group. Br is above I in group 17, so Br is more reactive.

# Brief Explanations:
- For nonmetals (especially halogens and group 16 nonmetals), reactivity decreases down a group because electronegativity decreases (harder to gain electrons).
# Answer:
F (more reactive than Cl: group 17, reactivity decreases down).
S (more reactive than Se: group 16, reactivity decreases down).
Br (more reactive than I: group 17, reactivity decreases down).

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QUESTION IMAGE

Question

Explanation:

Question 14: What is electronegativity?

Answer:

Question 15: How does electronegativity change across a period? Why?