test review: periodicity
3s²3p¹
4s²3d¹⁰4p⁵
periodic trends:
1. what is periodicity?
2. what is atomic radius?
3. how does the atomic radius change going across a period? why?
4. how does the atomic radius change going down a group? why?
5. for each of the following pairs, circle the larger atom.
a. y or rb
f. k or na
b. li or be
g. be or b
c. b or c
h. n or o
d. mg or sr
i. ba or ra
e. na or mg
j. ne or f
6. what is an ion? how is it different from a neutral atom?
7. what is the name of a positively charged ion? ______ does it gain or lose electrons? ______ are they usually metals or non - metals? ______
8. what is the name of a negatively charged ion? ______ does it gain or lose electrons? ______ are they usually metals or non - metals? ______
9. which is larger, k or k⁺¹? why?
10. which is larger, br, or br⁻¹? why?
11. ______ (metal or non - metal) ionic radii are larger than its atomic radius
12. ______ (metal or non - metal) ionic radii are smaller than its atomic radius.
13. for each of the following elements in each pair, decide whose ion (ionic radii) will be smaller
a. na or br
b. ca or f
c. k or c
d. o or mg

Question

test review: periodicity
3s²3p¹
4s²3d¹⁰4p⁵
periodic trends:
1. what is periodicity?
2. what is atomic radius?
3. how does the atomic radius change going across a period? why?
4. how does the atomic radius change going down a group? why?
5. for each of the following pairs, circle the larger atom.
   a. y or rb
   f. k or na
   b. li or be
   g. be or b
   c. b or c
   h. n or o
   d. mg or sr
   i. ba or ra
   e. na or mg
   j. ne or f
6. what is an ion? how is it different from a neutral atom?
7. what is the name of a positively charged ion? ______ does it gain or lose electrons? ______ are they usually metals or non - metals? ______
8. what is the name of a negatively charged ion? ______ does it gain or lose electrons? ______ are they usually metals or non - metals? ______
9. which is larger, k or k⁺¹? why?
10. which is larger, br, or br⁻¹? why?
11. ______ (metal or non - metal) ionic radii are larger than its atomic radius
12. ______ (metal or non - metal) ionic radii are smaller than its atomic radius.
13. for each of the following elements in each pair, decide whose ion (ionic radii) will be smaller
    a. na or br
    b. ca or f
    c. k or c
    d. o or mg

Accepted Answer

### Question 1: What is Periodicity?
# Brief Explanations:
Periodicity refers to the repeating pattern of chemical and physical properties of elements in the periodic table as their atomic number increases. This occurs because of the periodic nature of electron configurations, where elements with similar electron arrangements (in the same group) show similar properties. For example, alkali metals (Group 1) all have one valence electron and exhibit similar reactivity, like reacting vigorously with water.
# Answer:
Periodicity is the repeating pattern of chemical and physical properties of elements in the periodic table with increasing atomic number, due to periodic electron - configuration trends.

### Question 2: What is atomic radius?
# Brief Explanations:
Atomic radius is a measure of the size of an atom. It can be defined as the distance from the nucleus of an atom to the outermost shell of electrons. There are different types of atomic radii (covalent radius, metallic radius, van der Waals radius) depending on the bonding state of the atom. For a non - bonded atom, the van der Waals radius is used, while for a bonded atom (e.g., in a molecule), the covalent radius (for non - metals) or metallic radius (for metals) is considered.
# Answer:
Atomic radius is the distance from an atom's nucleus to its outermost electron shell, with different types (covalent, metallic, van der Waals) based on bonding.

### Question 3: How does the atomic radius change going across a period? WHY?
# Brief Explanations:
## Step1: Identify the trend across a period
As we move from left to right across a period (horizontal row) in the periodic table, the atomic radius generally decreases.
## Step2: Explain the reason
The number of protons in the nucleus (atomic number) increases as we move across a period. At the same time, electrons are added to the same valence shell. The increased positive charge of the nucleus exerts a stronger pull on the valence electrons, pulling them closer to the nucleus. Since the electrons are added to the same shell, there is no significant increase in electron - electron shielding to counteract this increased nuclear pull. For example, in the second period, from Li to Ne, the atomic number increases from 3 to 10, and the atomic radius decreases.
# Answer:
Across a period, atomic radius decreases. This is because the atomic number (protons) increases, increasing nuclear pull on valence electrons (added to the same shell, with no major shielding increase), pulling electrons closer to the nucleus.

### Question 4: How does the atomic radius change going down a group? WHY?
# Brief Explanations:
## Step1: Identify the trend down a group
When moving down a group (vertical column) in the periodic table, the atomic radius generally increases.
## Step2: Explain the reason
As we go down a group, the number of electron shells (energy levels) increases. Even though the number of protons in the nucleus also increases, the additional electron shells are further from the nucleus. The inner electron shells shield the valence electrons from the full positive charge of the nucleus. So, the valence electrons experience a weaker effective nuclear charge and are not pulled as close to the nucleus. For example, in Group 1 (alkali metals), from Li to Cs, the number of electron shells increases from 2 to 6, and the atomic radius increases.
# Answer:
Down a group, atomic radius increases. This is because the number of electron shells increases, and inner shells shield valence electrons from the nucleus, reducing the effective nuclear pull on valence electrons.

### Question 5: For each of the following pairs, circle the LARGER atom.
#### a. Y or Rb
# Brief Explanations:
Y (Yttrium) and Rb (Rubidium) are in the same period (period 5). As we move from left to right across a period, atomic radius decreases. Rb is to the left of Y in period 5. So, Rb has a larger atomic radius.
# Answer:
Rb

#### b. Li or Be
# Brief Explanations:
Li (Lithium) and Be (Beryllium) are in the same period (period 2). Moving from left to right across a period, atomic radius decreases. Li is to the left of Be. So, Li has a larger atomic radius.
# Answer:
Li

#### c. B or C
# Brief Explanations:
B (Boron) and C (Carbon) are in the same period (period 2). As we move from left to right across a period, atomic radius decreases. B is to the left of C. So, B has a larger atomic radius.
# Answer:
B

#### d. Mg or Sr
# Brief Explanations:
Mg (Magnesium) and Sr (Strontium) are in the same group (Group 2). Moving down a group, atomic radius increases. Sr is below Mg in Group 2. So, Sr has a larger atomic radius.
# Answer:
Sr

#### e. Na or Mg
# Brief Explanations:
Na (Sodium) and Mg (Magnesium) are in the same period (period 3). Moving from left to right across a period, atomic radius decreases. Na is to the left of Mg. So, Na has a larger atomic radius.
# Answer:
Na

#### f. K or Na
# Brief Explanations:
K (Potassium) and Na (Sodium) are in the same group (Group 1). Moving down a group, atomic radius increases. K is below Na in Group 1. So, K has a larger atomic radius.
# Answer:
K

#### g. Be or B
# Brief Explanations:
Be (Beryllium) and B (Boron) are in the same period (period 2). Moving from left to right across a period, atomic radius decreases. Be is to the left of B. So, Be has a larger atomic radius.
# Answer:
Be

#### h. N or O
# Brief Explanations:
N (Nitrogen) and O (Oxygen) are in the same period (period 2). Moving from left to right across a period, atomic radius decreases. N is to the left of O. So, N has a larger atomic radius.
# Answer:
N

#### i. Ba or Ra
# Brief Explanations:
Ba (Barium) and Ra (Radium) are in the same group (Group 2). Moving down a group, atomic radius increases. Ra is below Ba in Group 2. So, Ra has a larger atomic radius.
# Answer:
Ra

#### j. Ne or F
# Brief Explanations:
Ne (Neon) and F (Fluorine) are in the same period (period 2). Neon is a noble gas, and its atomic radius is measured as a van der Waals radius, while fluorine's atomic radius is a covalent radius. Also, as we move from left to right across a period, atomic radius generally decreases, but noble gases have larger radii due to the way their radii are measured (van der Waals radius is larger than covalent radius for adjacent elements). So, Ne has a larger atomic radius.
# Answer:
Ne

### Question 6: What is an ION? How is it different from a neutral atom?
# Brief Explanations:
An ion is an atom or a group of atoms that has a net electric charge. A neutral atom has an equal number of protons (positive charge) and electrons (negative charge), so its net charge is zero. An ion is formed when an atom gains or loses electrons. If an atom loses electrons, it becomes a positively charged ion (cation), and if it gains electrons, it becomes a negatively charged ion (anion). For example, a sodium atom (Na) has 11 protons and 11 electrons (neutral). When it loses one electron, it becomes a sodium ion ($Na^+$) with 11 protons and 10 electrons (net positive charge). A chlorine atom (Cl) has 17 protons and 17 electrons (neutral). When it gains one electron, it becomes a chloride ion ($Cl^-$) with 17 protons and 18 electrons (net negative charge).
# Answer:
An ion is a charged atom/group of atoms (gains/loses electrons). A neutral atom has equal protons and electrons (net charge 0); an ion has unequal protons and electrons (net charge ≠ 0).

### Question 7: What is the name of a positively charged ion? Does it gain or lose electrons? Are they usually metals or non - metals?
# Brief Explanations:
A positively charged ion is called a cation. To form a cation, an atom loses electrons. This is because losing electrons reduces the number of negative charges, so the positive charge from the protons dominates. Metals generally have low ionization energies, meaning it is relatively easy for them to lose electrons. So, cations are usually formed by metals. For example, sodium (a metal) loses an electron to form $Na^+$, and magnesium (a metal) loses two electrons to form $Mg^{2+}$.
# Answer:
Name: Cation; Loses electrons; Usually metals.

### Question 8: What is the name of a negatively charged ion? Does it gain or lose electrons? Are they usually metals or non - metals?
# Brief Explanations:
A negatively charged ion is called an anion. To form an anion, an atom gains electrons. Gaining electrons increases the number of negative charges, resulting in a net negative charge. Non - metals have relatively high electron affinities, meaning they have a tendency to gain electrons. So, anions are usually formed by non - metals. For example, chlorine (a non - metal) gains an electron to form $Cl^-$, and oxygen (a non - metal) gains two electrons to form $O^{2-}$.
# Answer:
Name: Anion; Gains electrons; Usually non - metals.

### Question 9: Which is larger, K or $K^+$? WHY?
# Brief Explanations:
A potassium atom (K) has an electron configuration of $1s^22s^22p^63s^23p^64s^1$. A potassium ion ($K^+$) is formed when K loses its 4s electron, so its electron configuration is $1s^22s^22p^63s^23p^6$. The $K^+$ ion has one less electron shell (the 4s shell is removed) compared to the K atom. Also, the remaining electrons in $K^+$ experience a greater effective nuclear charge because there are fewer electrons to shield the positive charge of the nucleus. So, the K atom is larger than the $K^+$ ion.
# Answer:
K is larger. K has an extra electron shell (4s) and more electrons, so less effective nuclear pull on its electrons compared to $K^+$ (which lost the 4s electron, has fewer electrons, and greater nuclear pull).

### Question 10: Which is larger, Br or $Br^-$? WHY?
# Brief Explanations:
A bromine atom (Br) has an electron configuration of $1s^22s^22p^63s^23p^63d^{10}4s^24p^5$. A bromide ion ($Br^-$) is formed when Br gains an electron, so its electron configuration is $1s^22s^22p^63s^23p^63d^{10}4s^24p^6$. When Br gains an electron, the number of electrons increases. This leads to increased electron - electron repulsion in the valence shell. The added electron also increases the shielding effect, reducing the effective nuclear charge on the valence electrons. As a result, the valence electrons are not pulled as close to the nucleus, and the radius of $Br^-$ is larger than that of Br.
# Answer:
$Br^-$ is larger. Br gains an electron to form $Br^-$, increasing electron - electron repulsion and shielding, reducing effective nuclear pull on electrons, so $Br^-$ has a larger radius.

### Question 11: (metal or non - metal) ionic radii are LARGER than its atomic radius
# Brief Explanations:
Non - metals form anions (negative ions) by gaining electrons. When a non - metal gains electrons, the number of electrons increases. This causes increased electron - electron repulsion in the valence shell, and the electrons are not pulled as close to the nucleus. Also, the added electrons increase the shielding effect, reducing the effective nuclear charge on the valence electrons. So, the ionic radius of a non - metal (anion) is larger than its atomic radius. For example, the atomic radius of Cl is smaller than the ionic radius of $Cl^-$.
# Answer:
Non - metal

### Question 12: (metal or non - metal) ionic radii are SMALLER than its atomic radius
# Brief Explanations:
Metals form cations (positive ions) by losing electrons. When a metal loses electrons, it loses an entire electron shell (in most cases). Also, the remaining electrons experience a greater effective nuclear charge because there are fewer electrons to shield the positive charge of the nucleus. This pulls the remaining electrons closer to the nucleus, resulting in a smaller ionic radius compared to the atomic radius. For example, the atomic radius of Na is larger than the ionic radius of $Na^+$.
# Answer:
Metal

### Question 13: For each of the following elements in each pair, decide whose ion (ionic radii) will be SMALLER
#### a. Na or Br
# Brief Explanations:
Sodium (Na) is a metal and forms a cation ($Na^+$) by losing one electron. Bromine (Br) is a non - metal and forms an anion ($Br^-$) by gaining one electron. Cations are generally smaller than anions because cations lose electron shells (or have a greater effective nuclear pull on fewer electrons) and anions gain electrons (increasing electron - electron repulsion and shielding). So, the ion of Na ($Na^+$) is smaller than the ion of Br ($Br^-$).
# Answer:
Na

#### b. Ca or F
# Brief Explanations:
Calcium (Ca) is a metal and forms a cation ($Ca^{2+}$) by losing two electrons. Fluorine (F) is a non - metal and forms an anion ($F^-$) by gaining one electron. Cations are smaller than anions. Also, the $Ca^{2+}$ ion has a relatively high positive charge and loses two electron shells (from the original Ca atom's electron configuration), while the $F^-$ ion gains an electron. The effective nuclear charge on $Ca^{2+}$ is relatively high, pulling the remaining electrons close. So, the ion of Ca ($Ca^{2+}$) is smaller than the ion of F ($F^-$).
# Answer:
Ca

#### c. K or C
# Brief Explanations:
Potassium (K) is a metal and forms a cation ($K^+$) by losing one electron. Carbon (C) is a non - metal and can form an anion ($C^{4 -}$) by gaining four electrons (or a cation, but in most cases, for ionic radius comparison here, we consider the common ionic states). The $K^+$ ion has a larger atomic number (more protons) and, although it loses an electron, the effective nuclear charge on its remaining electrons is significant. The $C^{4 -}$ ion gains four electrons, leading to a large increase in electron - electron repulsion. However, the $K^+$ ion has a smaller ionic radius because it is a cation (loses electrons, smaller size) and also has a higher effective nuclear charge compared to the $C^{4 -}$ ion. Wait, actually, Carbon usually forms covalent bonds, but if we consider ionic radius, the $K^+$ ion (radius around 133 pm) and the $C^{4 -}$ ion (if we consider it, but it's not a common stable ion in the same sense, but for the sake…

Sovi.AI - AI Math Tutor

QUESTION IMAGE

Question

Explanation:

Question 1: What is Periodicity?

Step1: Identify the trend across a period

Step2: Explain the reason

Answer:

Question 2: What is atomic radius?