AP Chemistry Unit 1 Study Guide Knowt

AP Chemistry Study Questions:

What is Avogadro’s number, and why is it important in chemistry?
How many molecules are in 2.00 moles of water (H₂O)?
What is the mass (in grams) of one mole of carbon dioxide (CO₂)?
If you have 18.0 grams of water (H₂O), how many moles do you have?
(Molar mass of H₂O = 18.0 g/mol)
How many atoms are present in 0.25 moles of sodium (Na)?
A sample contains 1.51 × 10²⁴ molecules of ammonia (NH₃). How many moles of ammonia are present?
The molar mass of sodium chloride (NaCl) is 58.5 g/mol. What is the mass of 0.50 moles of NaCl?
How many formula units are there in 5.00 grams of calcium carbonate (CaCO₃)?
(Molar mass of CaCO₃ = 100.1 g/mol)
Explain in your own words why chemists use the mole concept instead of counting individual atoms or molecules.
If a single molecule of glucose (C₆H₁₂O₆) has a mass of approximately 180 amu, what is the mass (in grams) of one mole of glucose?

AP Chemistry Study Questions: Mass Spectra and Isotopes

What information can you obtain about an element by analyzing its mass spectrum?
How does a mass spectrometer distinguish between different isotopes of the same element?
If a mass spectrum for an element shows two major peaks, what does this indicate about the element’s isotopes?
Why is the average atomic mass of an element usually not a whole number?
Describe how you would calculate the average atomic mass of an element using the mass spectrum data.
An element has two isotopes with masses 10.0 amu (abundance 20%) and 11.0 amu (abundance 80%). What is the average atomic mass?
What does the height or intensity of a peak in a mass spectrum represent?
Explain the relationship between the weighted average of isotopic masses and the periodic table value for atomic mass.
Why are samples containing only singly charged monatomic ions used for AP exam questions on mass spectra?
How could you use a mass spectrum to estimate the percent abundance of each isotope of an element?

AP Chemistry Study Questions: Empirical Formulas and Elemental Composition

Define empirical formula.
What is the law of definite proportions?
How does a molecular formula differ from an empirical formula?
A compound contains only carbon and hydrogen. If a 10.00 g sample contains 8.00 g C and 2.00 g H, what is the empirical formula?
Why are empirical formulas always written with the smallest whole number ratios?
Explain how to determine the percent by mass of each element in a compound given its formula.
A compound is 40.0% sulfur and 60.0% oxygen by mass. What is its empirical formula?
Given the formula unit NaCl, what does this tell you about the ratio of sodium to chlorine atoms?
A substance is found to contain 25.9% nitrogen and 74.1% oxygen by mass. What is the empirical formula?
If a compound’s empirical formula is CH2O and its molar mass is 180 g/mol, what is its molecular formula?
A 7.50 g sample of a compound contains 3.00 g magnesium and 4.50 g chlorine. Find the empirical formula.
What type of substances use formula units to represent their composition, and why?
For the compound K2O, what is the ratio of potassium atoms to oxygen atoms?
A sample contains 52.2% carbon, 13.0% hydrogen, and 34.8% oxygen. Find the empirical formula.
Why do ionic compounds not use molecular formulas?

AP Chemistry Study Guide Questions: Mixtures, Mass Composition, and Elemental Analysis (1.4.A)

What is the main difference between a pure substance and a mixture?
How can elemental analysis be used to determine if a sample is pure or a mixture?
A mixture contains 15.0 g of sodium chloride and 10.0 g of potassium chloride. What is the percent by mass of sodium chloride in the mixture?
How do the proportions of components in a mixture compare to those in a pure substance?
Describe a method you could use to determine the composition by mass of each component in a two-substance mixture.
A chemist analyzes a sample and finds it contains only one type of molecule. What does this indicate about the sample’s purity?
A mixture contains 5.0 g of compound X (molar mass = 50 g/mol) and 15.0 g of compound Y (molar mass = 30 g/mol). What is the mole fraction of compound X in the mixture?
Why can the composition of a mixture vary, but the composition of a pure substance cannot?
A mixture of calcium carbonate and magnesium carbonate has a total mass of 25.0 g. If it contains 12.0 g of calcium carbonate, what is the percent by mass of magnesium carbonate?
If elemental analysis of a sample shows it contains multiple elements but not in fixed ratios, what can you conclude about the sample?
Explain how you could use mass spectrometry to determine the purity of a sample.
A sample is found to be 60% by mass substance A and 40% by mass substance B. What does this say about the type of sample?
How does knowing the percent composition by mass of each component in a mixture help in chemical analysis?
If a mixture contains two substances, A and B, and the percent by mass of A is 25%, what is the percent by mass of B?
In a mixture, why is it important to know the mass of each component before performing a chemical reaction?

AP Chemistry Study Guide Questions: Electron Configuration and the Aufbau Principle (1.5.A)

What does the Aufbau principle state about the order in which electrons fill atomic orbitals?
Write the ground-state electron configuration for a neutral atom of carbon.
How many valence electrons does a neutral atom of chlorine have, and what is its electron configuration?
Describe the difference between core and valence electrons.
Give the ground-state electron configuration for the Na⁺ ion.
What is the electron configuration for a neutral atom of magnesium, and how does it change when it forms the Mg²⁺ ion?
Explain how the arrangement of electrons in shells and subshells is reflected in the periodic table.
Using the Aufbau principle, write the ground-state electron configuration for an atom of oxygen.
What is the electron configuration of Fe³⁺, and how does it compare to the neutral Fe atom?
According to Coulomb’s law, how does the force between two charged particles depend on the distance between them and their charges?

AP Chemistry Study Guide: Ionization Energy & Coulomb’s Law (1.5.A.4)

Define ionization energy.
According to Coulomb’s law, how does the distance between an electron and the nucleus affect the ionization energy?
Why is it generally harder to remove an electron from a 2s subshell than from a 3s subshell of the same atom?
Compare the ionization energy of Li and Na. Which is higher, and why?
Explain why the ionization energy increases as you move from left to right across a period.
Why does the ionization energy decrease as you move down a group in the periodic table?
Which would have a higher ionization energy: O or O⁺? Explain.
Why is the second ionization energy of an atom usually greater than the first?
How does effective nuclear charge (Zₑff) influence ionization energy?
Using Coulomb’s law, explain why it is more difficult to remove a core electron than a valence electron.

AP Chemistry Study Guide: Photoelectron Spectroscopy (PES) and Electron Structure (1.6.A)

What does a photoelectron spectrum (PES) measure?
How does each peak in a PES relate to electron configuration?
What information does the height of a peak in a PES provide?
What does the binding energy represented by a PES peak tell you about an electron’s distance from the nucleus?
Why do inner (core) electrons appear at higher binding energies in PES?
Sketch a possible PES for a lithium atom and label the peaks with their subshells.
How can you determine the number of electrons in a particular subshell from a PES?
Why do peaks for the 1s electrons of different elements appear at different binding energies?
If a PES for an atom shows three peaks, what might this indicate about its electron configuration?
How would the PES of Na⁺ differ from that of neutral Na?
Explain the relationship between the number of peaks and the number of occupied subshells.
How does increasing nuclear charge affect the position of PES peaks?
Why do valence electrons have lower binding energies than core electrons?
In PES, what does it mean if two peaks have the same height?
Describe how PES can be used to confirm the ground-state electron configuration of an atom.

AP Chemistry Study Guide: Atomic Properties, Trends, and Periodicity (1.7.A)

What is meant by periodicity in the context of the periodic table?
How does the ground-state electron configuration of an atom relate to its position on the periodic table?
Define effective nuclear charge.
How does shielding affect the attraction between the nucleus and valence electrons?
What is the general trend in atomic radius as you move across a period from left to right? Why?
Describe the trend in atomic radius as you move down a group. Why does this occur?
Define ionization energy.
What trend is observed for ionization energy as you move across a period? Explain why.
Describe the trend in ionization energy as you move down a group. Why does this happen?
What is electron affinity?
How does electron affinity generally change as you move from left to right across a period?
Define electronegativity.
Which element is the most electronegative, and why?
Explain the trend in electronegativity as you go down a group.
How does the presence of a completely filled shell affect an element’s reactivity?
Compare the size of a sodium atom to a sodium ion (Na⁺). Which is larger, and why?
Why do noble gases have very low or positive electron affinities?
Predict which has a higher first ionization energy: Mg or Al. Explain your reasoning.
Why does fluorine have a smaller atomic radius than oxygen?
How can periodic trends be used to predict the properties of an unknown element?

AP Chemistry Study Guide: Reactivity, Periodicity, and Chemical Bonding (1.8.A)

What determines whether two elements are likely to form a chemical bond?
How do valence electrons influence the reactivity of elements?
Why do elements in the same group of the periodic table often form similar compounds?
Give an example of two elements from the same group that form analogous compounds, and name the compounds.
Explain why alkali metals are so reactive.
Why do noble gases generally not react to form compounds?
How can you predict the typical ionic charge of an atom using the periodic table?
What is the typical charge for an ion formed by an element in Group 2?
Compare the reactivity of sodium and potassium, both alkali metals. Which is more reactive, and why?
Why do halogens readily form -1 ions?
Explain why the reactivity of the alkaline earth metals increases as you go down the group.
Describe how periodic trends help predict which elements will react together to form ionic compounds.
Which group of elements is known for forming colored compounds, and why?
How does the position of an element on the periodic table predict whether it will lose or gain electrons in a reaction?
Why do transition metals often form ions with different charges?
What type of compounds do elements with nearly full valence shells tend to form?
Explain why elements on the left side of the periodic table tend to be more reactive metals.
What is meant by "analogous compounds," and why are they important in chemistry?
How would you expect the reactivity of chlorine to compare with that of iodine? Explain.
Why do nonmetals, especially those in Group 17, tend to be very reactive?

Analysis (1.4.A) – Answer Key

1. What is the main difference between a pure substance and a mixture?
A pure substance contains only one type of particle (element or compound) with a fixed composition and definite properties. A mixture contains two or more substances physically combined, and the composition can vary.

2. How can elemental analysis be used to determine if a sample is pure or a mixture?
Elemental analysis determines the types and ratios of elements present. If results show a fixed, definite composition matching a known substance, the sample is pure. If multiple sets of ratios or different elements not in fixed proportion are found, the sample is a mixture.

3. A mixture contains 15.0 g of sodium chloride and 10.0 g of potassium chloride. What is the percent by mass of sodium chloride in the mixture?
Total mass = 15.0 g + 10.0 g = 25.0 g
% by mass NaCl = (15.0 g / 25.0 g) × 100 = 60%

4. How do the proportions of components in a mixture compare to those in a pure substance?
In a mixture, the proportions of components can vary. In a pure substance, the composition is fixed and does not change.

5. Describe a method you could use to determine the composition by mass of each component in a two-substance mixture.
One method is to physically separate the components (e.g., filtration, evaporation, chromatography), dry them, and weigh each component. The mass of each divided by the total gives percent by mass.

6. A chemist analyzes a sample and finds it contains only one type of molecule. What does this indicate about the sample’s purity?
This indicates the sample is a pure substance, since only one type of molecule is present.

7. A mixture contains 5.0 g of compound X (molar mass = 50 g/mol) and 15.0 g of compound Y (molar mass = 30 g/mol). What is the mole fraction of compound X in the mixture?
Moles X = 5.0 g / 50 g/mol = 0.10 mol
Moles Y = 15.0 g / 30 g/mol = 0.50 mol
Total moles = 0.10 + 0.50 = 0.60 mol
Mole fraction X = 0.10 / 0.60 = 0.167

8. Why can the composition of a mixture vary, but the composition of a pure substance cannot?
Mixtures are physical combinations, so their components can be present in any proportion. Pure substances have a definite, unchanging composition by definition.

9. A mixture of calcium carbonate and magnesium carbonate has a total mass of 25.0 g. If it contains 12.0 g of calcium carbonate, what is the percent by mass of magnesium carbonate?
Mass of MgCO₃ = 25.0 g – 12.0 g = 13.0 g
% by mass MgCO₃ = (13.0 g / 25.0 g) × 100 = 52%

10. If elemental analysis of a sample shows it contains multiple elements but not in fixed ratios, what can you conclude about the sample?
The sample is a mixture, not a pure compound.

11. Explain how you could use mass spectrometry to determine the purity of a sample.
Mass spectrometry can separate particles by mass. A pure substance will show a single peak (or predictable set of peaks for isotopes), while a mixture will show multiple peaks corresponding to different components.

12. A sample is found to be 60% by mass substance A and 40% by mass substance B. What does this say about the type of sample?
It is a mixture, since it contains two substances in a given proportion.

13. How does knowing the percent composition by mass of each component in a mixture help in chemical analysis?
It allows chemists to quantify the amount of each substance present, assess purity, and calculate how much of each component will react or be produced in a chemical process.

14. If a mixture contains two substances, A and B, and the percent by mass of A is 25%, what is the percent by mass of B?
100% – 25% = 75%

15. In a mixture, why is it important to know the mass of each component before performing a chemical reaction?
Knowing the mass of each component allows accurate calculation of reactant quantities, ensures the correct stoichiometry, and helps predict how much product will form or if a component will be in excess.

AP Chemistry Study Questions: Moles and Molar Mass

Avogadro’s number is 6.022 × 10²³ mol⁻¹. It represents the number of particles (atoms, molecules, or formula units) in one mole of a substance, allowing chemists to count particles by weighing samples.
Number of molecules = 2.00 mol × (6.022 × 10²³ molecules/mol) = 1.20 × 10²⁴ molecules.
The mass of one mole of CO₂ is the molar mass: 12.0 (C) + 2 × 16.0 (O) = 44.0 g.
n = m/M = 18.0 g / 18.0 g/mol = 1.00 mol.
Number of atoms = 0.25 mol × (6.022 × 10²³ atoms/mol) = 1.51 × 10²³ atoms.
n = (1.51 × 10²⁴ molecules) / (6.022 × 10²³ molecules/mol) = 2.51 mol.
Mass = n × M = 0.50 mol × 58.5 g/mol = 29.25 g.
**First, find moles: n = 5.00 g / 100.1 g/mol = 0.04995 mol
Number of formula units = 0.04995 mol × 6.022 × 10²³ = 3.01 × 10²² formula units.
**Counting individual atoms or molecules is impossible due to their extremely small size and huge numbers. The mole concept lets chemistsSomething went wrong. Please try again and contact support if the issue persists.

AP Chemistry Study Questions: Mass Spectra and Isotopes

It provides the identities and relative abundances of the isotopes present in a sample of the element.
Isotopes are distinguished by their different mass-to-charge ratios, resulting in separate peaks in the spectrum.
The element has at least two naturally occurring isotopes, each corresponding to a peak.
Because it is a weighted average of the masses of all naturally occurring isotopes, based on their relative abundances.
Multiply the mass of each isotope by its percent abundance (as a decimal), then sum the results for all isotopes.
(10.0 × 0.20) + (11.0 × 0.80) = 2.0 + 8.8 = 10.8 amu
The height or intensity of a peak shows the relative abundance of that isotope in the sample.
The periodic table lists the average atomic mass, which is the weighted average of all naturally occurring isotopic masses.
To simplify analysis and avoid complications from multiple charges or molecular ions, which are beyond the AP scope.
By comparing the relative heights (intensities) of the peaks, which represent the percent abundance of each isotope.

AP Chemistry Study Questions: Empirical Formulas and Elemental Composition

The empirical formula gives the simplest whole number ratio of atoms of each element in a compound.
The law of definite proportions states that a chemical compound always contains its component elements in a fixed ratio by mass.
The molecular formula shows the actual number of each type of atom in a molecule, while the empirical formula shows only the simplest ratio.
Find moles:
C: 8.00 g / 12.01 g/mol ≈ 0.667 mol
H: 2.00 g / 1.008 g/mol ≈ 1.98 mol
Divide by smallest:
C: 0.667/0.667 = 1 H: 1.98/0.667 ≈ 3
Empirical formula: CH₃
To reflect the most reduced form of the compound’s composition.
Divide the mass of each element in the compound by the total molar mass, then multiply by 100%.
S: 40.0 g / 32.07 g/mol ≈ 1.25 mol
O: 60.0 g / 16.00 g/mol ≈ 3.75 mol
Divide by smallest: S: 1, O: 3
Empirical formula: SO₃
There is a 1:1 ratio of sodium to chlorine atoms in NaCl.
N: 25.9 g / 14.01 g/mol ≈ 1.85 mol
O: 74.1 g / 16.00 g/mol ≈ 4.63 mol
Divide by smallest: N: 1, O: 2.5 Multiply by 2 for whole numbers: N:2, O:5
Empirical formula: N₂O₅
CH₂O molar mass = 30 g/mol. 180/30 = 6, so molecular formula = (CH₂O)₆ = C₆H₁₂O₆
Mg: 3.00 g / 24.31 g/mol ≈ 0.123 mol
Cl: 4.50 g / 35.45 g/mol ≈ 0.127 mol
Ratio ≈ 1:1 Empirical formula: MgCl
Ionic compounds, because they do not exist as discrete molecules but as a lattice of ions.
2 K : 1 O atom.
C: 52.2 g / 12.01 g/mol ≈ 4.35 mol
H: 13.0 g / 1.008 g/mol ≈ 12.9 mol
O: 34.8 g / 16.00 g/mol ≈ 2.18 mol
Divide by smallest (2.18):
C: 2 H: 6 O: 1
Empirical formula: C₂H₆O
Because ionic compounds do not consist of discrete molecules; their formulas represent the ratio of ions in the crystal lattice.

AP Chemistry Study Guide Questions: Electron Configuration and the Aufbau Principle (1.5.A)

What does the Aufbau principle state about the order in which electrons fill atomic orbitals?
Electrons fill the lowest energy orbitals first before occupying higher energy orbitals.
Write the ground-state electron configuration for a neutral atom of carbon.
1s² 2s² 2p²
How many valence electrons does a neutral atom of chlorine have, and what is its electron configuration?
Chlorine has 7 valence electrons. Its electron configuration is 1s² 2s² 2p⁶ 3s² 3p⁵.
Describe the difference between core and valence electrons.
Core electrons are the inner electrons that are not involved in bonding, while valence electrons are the outermost electrons that participate in chemical reactions.
Give the ground-state electron configuration for the Na⁺ ion.
1s² 2s² 2p⁶
What is the electron configuration for a neutral atom of magnesium, and how does it change when it forms the Mg²⁺ ion?
Neutral Mg: 1s² 2s² 2p⁶ 3s²
Mg²⁺ ion: 1s² 2s² 2p⁶ (loses the two 3s electrons)
Explain how the arrangement of electrons in shells and subshells is reflected in the periodic table.
Periods correspond to principal energy levels (shells), and blocks (s, p, d, f) correspond to subshells.
Using the Aufbau principle, write the ground-state electron configuration for an atom of oxygen.
1s² 2s² 2p⁴
What is the electron configuration of Fe³⁺, and how does it compare to the neutral Fe atom?
Neutral Fe: [Ar] 4s² 3d⁶
Fe³⁺: [Ar] 3d⁵ (loses two 4s electrons and one 3d electron)
According to Coulomb’s law, how does the force between two charged particles depend on the distance between them and their charges?
The force is directly proportional to the product of the charges and inversely proportional to the square of the distance between them (F ∝ q₁q₂/r²).

AP Chemistry Study Guide Questions: Electron Configuration and the Aufbau Principle (1.5.A.4)

Define ionization energy.
The energy required to remove the outermost electron from a gaseous atom or ion.
According to Coulomb’s law, how does the distance between an electron and the nucleus affect the ionization energy?
The greater the distance, the weaker the electrostatic attraction, so less energy is needed to remove the electron.
Why is it generally harder to remove an electron from a 2s subshell than from a 3s subshell of the same atom?
The 2s electron is closer to the nucleus and feels a stronger attraction, requiring more energy to remove.
Compare the ionization energy of Li and Na. Which is higher, and why?
Li has a higher ionization energy because its outer electron is closer to the nucleus than that of Na.
Explain why the ionization energy increases as you move from left to right across a period.
Nuclear charge increases while shielding remains similar, pulling electrons closer and making them harder to remove.
Why does the ionization energy decrease as you move down a group in the periodic table?
Outer electrons are farther from the nucleus and more shielded by inner electrons, so they are easier to remove.
Which would have a higher ionization energy: O or O⁺? Explain.
O⁺ has a higher ionization energy, since removing an electron from a positively charged ion requires more energy.
Why is the second ionization energy of an atom usually greater than the first?
After removing the first electron, the remaining electrons experience a stronger effective nuclear charge, making the next electron harder to remove.
How does effective nuclear charge (Zₑff) influence ionization energy?
Greater effective nuclear charge increases the attraction to electrons, raising the ionization energy.
Using Coulomb’s law, explain why it is more difficult to remove a core electron than a valence electron.
Core electrons are much closer to the nucleus and experience a much stronger electrostatic attraction, so more energy is needed to remove them.

AP Chemistry Study Guide: Photoelectron Spectroscopy (PES) and Electron Structure (1.6.A)

What does a photoelectron spectrum (PES) measure?
The energies required to remove electrons from the various subshells of an atom.
How does each peak in a PES relate to electron configuration?
Each peak corresponds to electrons in a specific subshell (e.g., 1s, 2s, 2p).
What information does the height of a peak in a PES provide?
The height is proportional to the number of electrons in that subshell.
What does the binding energy represented by a PES peak tell you about an electron’s distance from the nucleus?
Higher binding energy means the electron is closer to the nucleus.
Why do inner (core) electrons appear at higher binding energies in PES?
They are more strongly attracted to the nucleus and require more energy to remove.
Sketch a possible PES for a lithium atom and label the peaks with their subshells.
(Description: Two peaks—one larger (2 electrons) at higher energy for 1s, one smaller (1 electron) at lower energy for 2s.)
How can you determine the number of electrons in a particular subshell from a PES?
By comparing the relative heights of the peaks; higher peaks mean more electrons.
Why do peaks for the 1s electrons of different elements appear at different binding energies?
Because the nuclear charge changes, increasing the attraction for electrons in heavier elements.
If a PES for an atom shows three peaks, what might this indicate about its electron configuration?
The atom has electrons in three different subshells.
How would the PES of Na⁺ differ from that of neutral Na?
Na⁺ would have one fewer peak in the outermost subshell (since it lost an electron).
Explain the relationship between the number of peaks and the number of occupied subshells.
Each peak represents a different occupied subshell.
How does increasing nuclear charge affect the position of PES peaks?
Peaks shift to higher binding energies as nuclear charge increases.
Why do valence electrons have lower binding energies than core electrons?
They are further from the nucleus and less strongly attracted.
In PES, what does it mean if two peaks have the same height?
There are the same number of electrons in those two subshells.
Describe how PES can be used to confirm the ground-state electron configuration of an atom.
By matching the number and heights of peaks to the predicted number of electrons in each subshell.

AP Chemistry Study Guide: Atomic Properties, Trends, and Periodicity (1.7.A)

What is meant by periodicity in the context of the periodic table?
Periodicity refers to the recurring trends or patterns in the properties of elements as you move across periods or down groups in the periodic table.
How does the ground-state electron configuration of an atom relate to its position on the periodic table?
The electron configuration determines the element’s column (group) and row (period), reflecting its chemical reactivity and properties.
Define effective nuclear charge.
The net positive charge experienced by valence electrons, accounting for the shielding effect of inner electrons.
How does shielding affect the attraction between the nucleus and valence electrons?
Shielding by inner electrons reduces the attraction between the nucleus and outer (valence) electrons.
What is the general trend in atomic radius as you move across a period from left to right? Why?
Atomic radius decreases because the increasing nuclear charge pulls electrons closer.
Describe the trend in atomic radius as you move down a group. Why does this occur?
Atomic radius increases due to the addition of energy levels (shells), increasing the distance from the nucleus.
Define ionization energy.
The energy required to remove an electron from a gaseous atom or ion.
What trend is observed for ionization energy as you move across a period? Explain why.
Ionization energy increases because effective nuclear charge increases, making it harder to remove an electron.
Describe the trend in ionization energy as you move down a group. Why does this happen?
Ionization energy decreases due to greater distance from the nucleus and increased shielding.
What is electron affinity?
The energy change that occurs when an electron is added to a neutral atom in the gas phase.
How does electron affinity generally change as you move from left to right across a period?
It generally becomes more negative (releases more energy), as atoms more readily accept electrons.
Define electronegativity.
The tendency of an atom to attract electrons in a chemical bond.
Which element is the most electronegative, and why?
Fluorine, due to its high effective nuclear charge and small radius.
Explain the trend in electronegativity as you go down a group.
Electronegativity decreases because atoms are larger and the nucleus is less able to attract bonding electrons.
How does the presence of a completely filled shell affect an element’s reactivity?
Elements with filled shells (like noble gases) are very stable and chemically unreactive.
Compare the size of a sodium atom to a sodium ion (Na⁺). Which is larger, and why?
The sodium atom is larger; Na⁺ has lost an electron and an entire energy level, reducing its radius.
Why do noble gases have very low or positive electron affinities?
They have filled valence shells and gain no stability by adding another electron.
Predict which has a higher first ionization energy: Mg or Al. Explain your reasoning.
Mg has a higher first ionization energy; removing an electron from Mg disturbs a full 3s subshell, while Al starts filling the 3p subshell.
Why does fluorine have a smaller atomic radius than oxygen?
Fluorine has a higher nuclear charge, pulling electrons closer.
How can periodic trends be used to predict the properties of an unknown element?
By comparing the element’s position to known trends, predictions can be made about its atomic radius, ionization energy, electron affinity, and electronegativity.

AP Chemistry Study Guide: Reactivity, Periodicity, and Chemical Bonding (1.8.A)

The interactions between the valence electrons and nuclei of the elements determine if they will form a bond.
Valence electrons determine how easily an element can gain, lose, or share electrons, directly affecting reactivity.
Because they have the same number of valence electrons, leading to similar chemical behaviors.
Sodium and potassium (both Group 1) form NaCl and KCl, respectively—both are alkali metal chlorides.
Alkali metals have one valence electron that is easily lost, making them highly reactive.
Noble gases have full valence shells, so they are stable and rarely form compounds.
By counting how far the element is from a noble gas; main group metals lose electrons to achieve a noble gas configuration, nonmetals gain.
+2, because Group 2 elements lose two electrons to achieve a full shell.
Potassium is more reactive than sodium; its outer electron is farther from the nucleus and more easily lost.
Halogens have seven valence electrons and gain one more to complete their octet, forming -1 ions.
The outer electrons are farther from the nucleus and more easily lost, increasing reactivity down the group.
Elements on the left (metals) tend to lose electrons, elements on the right (nonmetals) tend to gain; opposites attract to form ionic compounds.
Transition metals, due to their partially filled d subshells allowing various electron transitions.
Elements on the left lose electrons (form cations); elements on the right gain them (form anions).
Their d subshells allow for multiple stable oxidation states.
They tend to form anions, often in molecular or ionic compounds.
They have few valence electrons and lose them easily, making them highly reactive.
Compounds with similar formulas and structures formed by elements in the same group; important for recognizing chemical patterns.
Chlorine is more reactive because it is higher in Group 17, with a greater tendency to gain an electron.
They are close to completing their valence shell and readily gain electrons, making them highly reactive.