unit 1: atomic structure and properties

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Unit 1: Atomic Structure and Properties

www.leffellabs.com Schedule

A B Content:

9/15 9/17 Moles and Molar Mass 9/20 9/21 Mass Spectroscopy, Composition of Pure Substances

9/22 9/23 Composition of Mixtures

9/24 9/27 Lab 1

9/28 9/29 Atomic Structure

9/30 10/1 Photoelectron Spectroscopy

10/4 10/5 Periodic Trends, Valence Electrons and Ionic Charge 10/6 10/7 Mastery Check 1

Graded Assignments and Due Dates

• OWL Assignment 1, due 9/26

• OWL Assignment 2, due 10/3

• Lab 1 Purification of a Mixture, due 10/4 (A) and 10/5 (B)

• Unit 1 Focused Notes, due 10/6 (A) and 10/7 (B)

• Unit 1 Progress Check Reflection, due 10/6 (A) and 10/7 (B) Recommended Readings

• Section 1.5 Significant Figures in Calculations

• Section 1.7 Dimensional Analysis

• Section 3.1 Counting by Weighing

• Section 3.2 Atomic Masses

• Section 3.3 The Mole

• Section 3.4 Molar Mass

• Section 3.6 Percent Composition of Compounds

• Section 3.7 Determining the Formula of a Compound

• Section 7.12 Periodic Trends in Atomic Properties

• Section 8.4 Ions: Electron Configurations and Sizes Practice/ Review Questions

• Chapter 3 Review Questions, #1-10, pp.112-113

• Chapter 3 Exercises #39, 43, 73, 77, 89, pp. 115-115c

• Chapter 7 Exercises #115, 119, 121, 133, pp.299d-299f

• Chapter 8 Exercises #51, 55 p. 353b

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1.1 Moles and Molar Mass

Outcome: I can calculate quantities of a substance or its relative number of particles using dimensional analysis and the mole concept.

Counting by Weighing

1. Watch the video Surprising my Dog with 1,000,000 Pieces of Dog Food. (https://bit.ly/3BBlFEU) Do “The Math,” showing work. Use metric units.

The Mole

2. A formal definition for the mole (mol) is the number of carbon atoms in exactly 12 grams of pure carbon-12. How does this relate to the dog food example?

3. The sample of pure carbon-12 described above contains 6.02×1023 atoms. Express this number as a conversion factor between moles and atoms.

https://youtu.be/CGlTe0jDtfg?t=45

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Molar Mass The molar mass of a substance the mass in grams of 1 mole of the compound, with units of grams per mol. The molar mass of methane, CH4, is calculated like this:

Mass of 1 mole C =12.01 g

Mass of 4 mol H =4×1.008 g

Mass of 1 mol CH4 =16.04 g

Methane is a molecular compound, so we count by the number of molecules (1 mol CH4 = 6.02×1023 molecules). Because ionic compounds do not contain molecules, we use the term formula unit as the basic unit for these materials. For example, the formula unit for table salt is NaCl, which contains 1 mol of Na+ ions and 1 mol of Cl− ions and has a molar mass of 58.44 g/mol (22.99 g/mol + 35.45 g/mol).

4. Compare the number of atoms present in 12.01 grams of carbon to the number of atoms present in 26.98 grams of aluminum.

1. Rank three samples in order of increasing number of particles, increasing mass, and increasing mole amounts:

Sample A: 1.0 mole of carbon Sample B: 18 grams of carbon monoxide Sample C: 3.0×1023 molecules of water)

Reflection Questions: 1. It is not possible to count the number of atoms or molecules in a sample. How does

the mass of substance connect to the number of particles present? 2. How does Avogadro’s number relate to the number of particles in a sample? 3. How does the mole allow different quantities to be compared?

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1.2 Mass Spectroscopy

Outcome: I can explain the quantitative relationship between the mass spectrum of an element and the masses of the element’s isotopes.

Isotopes Review

1. Does the sample contain identical atoms of copper? Explain.

2. What percentage of the copper atoms are 63Cu? What is the percent of 65Cu?

3. Use a periodic table to answer the following: a. What is the average atomic mass of copper?

b. Is this number closer to 63 amu or 65 amu?

c. How does the information given in the picture relate to Question 3b?

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What Happens in a Mass Spectrometer? Mass spectrometry was developed during the search for isotopes almost 100 years ago and has evolved into a powerful tool for analyzing the mass and structure of compounds. It is widely used in laboratories across the world for forensic analysis of trace amounts of substances, and in research to determine the structures of large and complex molecules, such as proteins and nucleic acids. As AP chemistry students, you’ll need to understand mass spectra as they relate to isotopes - now is a good time to review the outcome.

• A sample is injected into the machine, which turns it into a vapor. The vapor passes into an ionization chamber, where a stream of electrons hit the atoms of the sample, knocking electrons off from the atoms. This gives the atoms a positive charge, forming cations.

• Next, cations are accelerated so that they all travel at the same speed. They then enter a magnetic field, which deflects the cations based on their mass. Lighter cations are deflected more than heavier ones.

• Finally, the cations strike a detector. This causes an electric current, which is then interpreted by a computer to produce the mass spectrum of the sample.

• This is all done under a vacuum, to avoid the cations from colliding with (and possibly reacting with) molecules in air.

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4. Watch Mass Spectroscopy by Bozeman Science (https://bit.ly/2WJD13P) and answer the following.

a. What is the purpose of a mass spectrometer?

b. Which statements of Dalton’s atomic theory are inaccurate? Explain why.

c. Describe the main components of a mass spectrometer. How does the sample move through the machine to reach the detector?

d. Imagine a cannonball is about to travel past you, and you want to deflect its path using a jet of water from a hose. The cannonball is so heavy that the water jet doesn’t do much. If it were a tennis ball, the water jet would cause a lot of deflection. How does this metaphor relate to the functioning of a mass spectrometer?

e. What is the relationship between the number and relative height of the peaks and the isotopes of an element?

https://www.youtube.com/watch?v=mBT73Pesiog

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5. The graph was produced when lithium was analyzed in a mass spectrometer.

a. How many isotopes of lithium exist?

b. What is the mass of the most abundant isotope? The least abundant isotope?

c. Without performing any calculations, predict the approximate average atomic mass of lithium. Explain the basis for your prediction.

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6. The height of each peak is the relative intensity, not the percent abundance. You will need to calculate the percent abundance then the average atomic mass.

a. What is the relative intensity of each peak?

b. What is the total intensity of the peaks? Find this by adding them.

c. What is the percent of the intensity of each peak?

% =𝑝𝑎𝑟𝑡

𝑤ℎ𝑜𝑙𝑒× 100

d. You’ve just found the percent abundance using the mass spectrum. Complete the table below and calculate the average atomic mass.

Isotope Isotope Mass % Abundance

Lithium-6 6 amu

Lithium-7 7 amu

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7. Consider the element boron (atomic number 5). a. What is the average atomic mass of boron?

b. If there are two predominant isotopes of boron, 10B and 11B, which do you think is the most abundant?

c. Estimate the percentage of each isotope of boron.

d. Draw a sketch for the mass spectrum of boron. Label each axis and label each peak with the isotope it represents.

Reflection Questions: 1. How does a mass spectrometer separate apart atoms? Explain. 2. Explain how a mass spectrum of an element can be used to identify isotopes and the

relative abundance of the isotope. 3. How is mass spectroscopy used to determine the average atomic mass of an element?

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1.3 Composition of Pure Substances

Outcome: I can explain the quantitative relationship between the elemental composition by mass and the empirical formula of a pure substance.

Empirical Formula A common method for determining the chemical formula of a hydrocarbon involves burning it with excess oxygen. All the carbon is converted to CO2, while all the hydrogen is converted into H2O. In the diagram below, the mass of the absorbing chambers was measured before and after the reaction, allowing us to measure the amount of H2O and CO2 produced. The masses of these compounds can be used to find the percent composition of the hydrocarbon.

1. A 0.1156 g sample of a substance made from carbon, hydrogen, and nitrogen was burned in excess oxygen, producing 0.1638 g of CO2 and 0.1676 g of H2O.

a. How many grams of carbon are in 0.1638 g of CO2? (HINT: use a mole ratio of C to CO2).

b. How many grams of hydrogen are in 0.1676 g of H2O?

c. Find the percent by mass of carbon in the original sample.

d. Find the percent by mass of hydrogen in the original sample.

e. What is the percent by mass of nitrogen? Remember the substance is only made from carbon, hydrogen, and nitrogen.

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2. What if during the experiment, the flow of oxygen is interrupted, resulting in some of the carbon forming CO instead of CO2. How would this impact the percent compositions you found?

3. In the previous question, your answers should be 38.67% C, 16.22% H and 45.11% N. Use these values to calculate the empirical formula of the compound.

The empirical formula is the simplest whole number ratio of atoms in a molecule. The molecular formula will be some multiple of the empirical formula. For example, the formula for glucose is six times the empirical formula (CH2O)6.

C6H12O6 = (CH2O)6 To specify the exact formula of the molecule, we need to see how many times the empirical molar mass goes into the molar mass of the compound.

4. The correct formula from Question 2 is CH5N. a. What is the formula mass of CH5N?

b. If the molar mass of the compound is 62.1 g/mol, what is the molecular formula?

Reflection Questions: 1. The law of definite proportions says that the percent by mass of each element in a

compound is constant. How does this relate to the empirical formula of a compound? 2. Without doing any calculations, which has a greater percent by mass of hydrogen, H2O

or H2O2? Explain.

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1.4 Composition of Mixtures

Outcome: I can explain the quantitative relationship between the elemental composition by mass and the composition of substances in a mixture.

Chemical Analysis Chemical analysis is a useful tool for determining the formula of a compound, but also for determining the composition of a mixture. One useful way to determine the purity of a sample is to compare the percent composition of the elements in the compound to known values.

1. A sample of sodium chloride is believed to be contaminated. Upon analysis, the sample was found to be 55% chlorine by mass. Which of the following salts could be the contaminant? Justify your answer using math.

NaI KCl LiCl

2. A sample of sodium carbonate (Na2CO3) has been contaminated with sodium hydrogen carbonate (NaHCO3). Are there physical properties that are different for the two substances that would allow us to separate them?

3. When heated, sodium hydrogen carbonate will decompose into sodium carbonate. Sodium carbonate does not decompose when heated.

2NaHCO3(s) + heat → Na2CO3(s) + H2O(g) + CO2(g) Na2CO3(s) + heat → Na2CO3(s) (no reaction)

a. Which graph shows how the mass of the NaHCO3 will change during the reaction?

b. Which graph shows how the mass of the contaminated mixture will change during the reaction?

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4. The balanced equation for the experiment is listed below. The desired product is solid

metal carbonate. Calculate the atom economy (https://bit.ly/3BDFcob) for the reaction. 2NaHCO3(s) → Na2CO3(s) + H2O(g) + CO2(g)

5. Review the stoichiometry of the bicarbonate decomposition. How does the mass that is lost relate to the starting mass of material?

2NaHCO3(s) → Na2CO3(s) + H2O(g) + CO2(g)

6. Review the Principles of Green Chemistry (https://bit.ly/3tdjoN4). What is a possible consideration from the principles of green chemistry that could help you to compare the efficiency or “greenness” of two reactions?

Reflection Questions: 1. Compare pure substances and mixtures in terms of composition (types of molecules or

formula units involved). 2. How can percent composition be used as a measure of the purity of a sample? Why

might chemists need to know the purity of a chemical?

https://www.youtube.com/watch?v=Y9NVz7_9zw4

https://www.acs.org/content/acs/en/greenchemistry/principles/12-principles-of-green-chemistry.html

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1.5 Atomic Structure

Outcome: I can represent the electron configuration of an element or ions of an element using the Aufbau principle.

The Boarding House

1. Examine the boarding house diagrams. Match the symbol with its meaning.

_____ _____1s2 2s2 2p6 3s1

_____

A. Bunk bed for boarders B. Manager’s code for occupied rooms C. Boarder

2. Look at the manager’s code for each diagram. Using the manager’s code:

1s2 2s2 2p6 3s1 a. Underline the floor numbers b. Circle the types of rooms c. Highlight the numbers of boarders in each room

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3. Consider the statements below and determine what best describes the manager’s set of rules for placing boarders in bunks and rooms.

a. The boarding house will rent out beds on the ______ floor first. i. 1st

ii. 2nd iii. 3rd

b. Boarders are only allowed to double up in a bunk in a room when _____.

i. there is an even number of boarders ii. all bottom bunks are filled

c. The next floor will be opened when _____ on the floor below are occupied.

i. half the bunks ii. at least one of the rooms

iii. all the bunks

d. The pink room on a floor will be opened for boarders when _____. i. all lower bunks in the sunny room on that floor are occupied

ii. all the bunks in the sunny room on that floor are occupied iii. the sunny room on that floor is open

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Ground State Orbital Diagrams and Electron Configurations

4. Examine the orbital diagrams and electron configurations shown. Match the symbol

with its meaning.

_____

_____

_____

_____

_____

A. Single electron B. Pair of electrons with opposite spins C. Atomic orbital (region of space around the

nucleus where an electron is likely to be found) D. Sublevel (set of orbitals with equivalent

energy) E. Electron configuration

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5. Examine the orbital diagrams and electron configurations. Using the electron configuration:

1s2 2s2 2p4 a. Underline the energy levels. b. Circle the sublevels. c. Highlight the numbers of electrons.

6. The lowest potential energy arrangement of electrons in an atom is called the ground

state. Ground state electron configurations can be predicted by a strict set of rules known as the Aufbau principle (“Aufbau” means building up). Examine the diagrams and the statements below to determine the phrase that best describes each rule.

a. Based on where a single electron is placed, the lowest potential energy electron in an atom is found in the _____ sublevel.

i. 1s ii. 2s

iii. 3s

b. Electrons will occupy a p-orbital only after _____ . i. the previous s-orbital is half full

ii. the previous s-orbital is completely full iii. the previous s-orbital is empty

c. Electrons can begin to occupy energy levels with the next highest integer

designation (e.g.,2 vs. 1, 3 vs. 2) only after _____ on the energy level below it are occupied.

i. half of the orbitals ii. at least one of the orbitals

iii. all the orbitals

7. The Pauli exclusion principle describes the restriction on the placement of electrons into the same orbital. The Pauli exclusion principle can be expressed as: “If two electrons occupy the same orbital, they must have _____.”

i. the same spin ii. opposite spins

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8. Hund’s rule describes how electrons are distributed among orbitals of the same sublevel when there is more than one way to distribute them.

a. Electrons will pair up in an orbital only when _____. i. there is an even number of electrons in the sublevel

ii. all orbitals in the same sublevel have one electron b. When single electrons occupy different orbitals of the same sublevel, ____.

i. they all have the same spin ii. they all have different spins

iii. their spins are random

9. For each of the symbols below, provide the name or description of the analogous component that was used in the boarding house model.

10. How could the boarding house model be modified to better represent the relative

energies of s and p sublevels?

11. Complete the ground state orbital energy level diagrams and write the electron configuration for each atom or ion.

Sulfur

Fluoride (F−)

Sodium ion (Na+)

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Excited State Electron Configuration An excited state electron configuration is any electron configuration for an atom that contains the correct total number of electrons but has a higher total electron potential energy than the ground state electron configuration.

12. Consider the orbital diagram below.

a. How many electrons are there in one atom of element X?

b. Identify element X and provide its ground state electron configuration. Assume that the atom is neutral.

c. Is the arrangement of electrons in the orbital diagram in Model 3 higher in total potential energy or lower in total potential energy than the ground state electron configuration of element X? Explain your reasoning.

Reflection Question: 1. Summarize how the Aufbau principle, Hund’s rule, and the Pauli exclusion principle tell

you the order to fill electrons into orbitals.

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1.6 Photoelectron Spectroscopy

Outcome: I can explain the relationship between the photoelectron spectrum of an atom or ion and:

a. The electron configuration of the atom b. The interactions between electrons and the nucleus.

Distance and Coulombic Attraction

1. The diagram above represents the force of attraction between a proton and an electron.

Based on your observations: a. If the distance between a proton and an electron is 0.60 nm, would the force of

attraction be greater than or less than 0.26×10−8? Why?

b. If two protons are 0.10 nm away from an electron, would the force of attraction be greater than or less than 2.30×10−8? Why?

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Number of Protons and Coulombic Attraction

2. The diagram above represents the force of attraction between several protons

(representative of nuclei) and an electron. Based on your observations: What would be the force of attraction between 5 protons and an electron that are 0.10 nm apart? Show math work to support your answer.

3. Using the data above, make a prediction about the force of attraction experienced by each electron in the diagram below.

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Potential Energy in an Atom

4. The diagram above shows two charged particles and the attraction between them. If q

for an electron is –1, a. What is q for a proton?

b. What is q for a neutron?

c. What is q for the nucleus of a nitrogen atom?

5. As distance increases, what happens to the value of the potential energy? Does the equation for potential energy make sense given your answer to Question 1? Explain.

6. Review your answer to Question 3. The attractive force experienced by each electron is about 4.60×10−8 N. This force is NOT “divided up” amongst the electrons.

a. Use the potential energy equation to explain why each electron feels the same force of attraction to the nucleus

b. Based on your last answer, are the repulsive forces between electrons significant (compared to the attractive forces)? Do you think this could change when several electrons are involved?

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Potential Energy and Ionization Energy

q1 q2 Distance (pm) IE (J) V (J)

Proton Electron 5000. 4.62 × 10–20

Proton Electron 1000. 2.31 × 10–19

Proton Electron 200.0 1.16 × 10–18

7. Do you expect the potential energy, V, of the atoms in the table above to be positive or

negative numbers? Explain.

8. Calculate the potential energies (V) for each atom. Use the value k = 2.31x10–16 J∙pm.

9. Ionization energy is the amount of energy required to pull an electron away from a nucleus. What is the relationship between ionization energy and potential energy? Specifically, why is one negative and one positive?

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Photoelectron Spectroscopy Photoelectron spectroscopy (PES) is used to study the ionization energies of an atom. To complete PES, the following steps occur:

• Photons (packets of energy) of equal energy are shot at an element in the gas phase.

• When a photon collides with an electron in an atom, it transfers its energy to the electron, knocking an electron out of the atom.

• If the energy of the photon is larger than the required ionization energy, the extra energy is converted into the kinetic energy of the ejected electron.

• A sensor detects the velocity of the ejected electron, allowing the kinetic energy of the electron to be found.

𝐾𝐸 =1

2𝑚𝑣2

• The energy of the photon is equal to the ionization energy plus the kinetic energy. If you know the energy of the photon and the kinetic energy of the ejected electron, you can determine the ionization energy:

𝐼𝐸 = 𝐸𝑝ℎ𝑜𝑡𝑜𝑛 − 𝐾𝐸

10. What is the ionization energy of the atom depicted below?

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Atoms contain several electrons arranged into energy levels. When a photon of high enough energy strikes an atom, any one of these electrons may be ejected, with equal probability. Thus, for a large group of identical atoms, electrons will come from all possible energy levels of the atom. This means that only a few different energies of ejected electrons will be obtained, corresponding to the energy levels in the atom. A photoelectron spectrum (PES) shows the number of ionized electrons versus their ionization energies.

11. Based on the diagram and PES above: a. What is the ionization energy of the ejected electron?

b. What determines the position of the peak on the horizontal axis?

c. How many energy levels (peaks) are present?

d. What determines the height (or intensity) of each peak?

e. A student claims the element shown in the PES is hydrogen, while another student claims the element shown is helium. Based on the number of energy levels, explain how both can be right.

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12. The atom depicted in the energy level diagram below has a total of five electrons, distributed amongst two energy levels.

a. Assuming a neutral atom, what element is this?

b. Complete the shell model of the atom above.

c. Suppose that the two energy levels are –0.85 MJ/mole and –4.25 MJ/mole. On the graph below, sketch the PES.

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13. Considering the diagram below,

a. What element is shown?

b. Draw the expected PES of this element. Indicate the relative intensity (peak size) and the position of the peak(s).

c. Why is the distance of the shell from the nucleus important when determining the peak position on a photoelectron spectrum?

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14. The actual PES for neon is above, along with the electron configuration for neon. There are three peaks in the spectrum for neon.

a. What is the reasoning behind the assumption that the peak at 84.0 MJ/mole corresponds to the electrons in the n = 1 shell?

b. Which orbital is closer to the nucleus: the 2s or 2p? Explain.

c. The ionization energies for peak 2 and 3 are very similar, while the value for peak 1 is much larger. Explain in terms of electron configuration.

Reflection Questions 1. What information can we get from the number of peaks in a PES? What about the

relative height of the peaks in a PES? 2. Why is it unimportant to place numbers along the vertical axis of PES? 3. How does the position of the peaks on a PES relate to the interactions between the

electrons and the nucleus?

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1.7 Periodic Trends

Outcome: I can explain the relationship between trends in atomic properties of elements and electronic structure and periodicity.

The Shell Model Electrons in atoms are attracted to the nucleus by a Coulombic force. Thus, energy must be supplied to pull the electron away, creating a positively charged cation, and a free electron.

M(g) → M+ + e– The ionization energy (IE) of an element is the minimum energy required to remove an electron from a gaseous atom of that element. Chemists report ionization energies per mole. For a H atom, IE = 2.178×10–18 J

1. How much total energy would it take to remove the electrons from a mole of H atoms? Write this energy in MJ/mole (1 MJ = 106 J).

2. For atoms with many electrons, not all electrons are at the same distance from the nucleus. In this case, which electron would have the lowest ionization energy: the electron that is closest to the nucleus or the electron that is farthest from the nucleus? Explain.

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3. The models below depict a hydrogen atom and a helium atom. Use the diagram to explain why the IE1 of He is approximately twice the IE1 of H.

4. Consider the two models of a Li atom below.

a. The IE of H is 1.31 MJ/mole. If all three electrons in Li were in the first shell

(model 1), which of the following values would be the better estimate of the IE1 for Li: 3.6 MJ/mole or 0.6 MJ/mole? Explain.

b. Examine the first ionization energy graph on the previous page. Notice that the value for Li is much lower than the value for H and He. How is this finding consistent with model 2?

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The electrons in the outermost shell of an atom are valence electrons, while electrons in shells closer to the nucleus are called inner-shell electrons. The nucleus plus the inner shells of electrons constitute the core of the atom, and the net overall charge on the core is called the core charge.

5. Considering the models for the beryllium atom above:

a. How many inner shell and valence electrons does Be have?

b. Show how the core charge for Be was calculated.

c. What is the relationship between the number of valence electrons and the core charge of the neutral atom?

6. If the valence shells of Li and Be are the same distance from their nuclei, explain how the core charges are consistent with the IE1 for Li (0.52 MJ/mol) and Be (0.9 MJ/mol).

7. Consider the core charge of Na and Li. a. How does the core charge of Na compare to the core charge of Li?

b. The IE1 for Na is 0.50 MJ/mol and the IE1 for Li is 0.52 MJ/mol. Based on this data, which atom has a smaller valence shell radius (which e– is closer to the nucleus)?

c. Neon has a higher IE1 value than Na. How is this consistent with their core charges?

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Atomic Size

Element Shell Core Charge Radius (pm)

B n = 2 +3 89

C n = 2 +4 77

O n = 2 +6 66

S n = 3 +6 104

As n = 4 +5 121

Se n = 4 +6 117

8. The data for several atoms is given. a. What is the relationship between the valence shell of each atom in the atomic

size table above and its position in the periodic table?

b. Why does the core charge increase as one moves from left to right across a period in the periodic table (for example, from boron to carbon to oxygen)?

c. What trend in atomic radius is observed as one moves from left to right across a period? Explain why this trend exists.

d. What trend in atomic radius is observed as one moves down a group in the periodic table? Explain why this trend exists.

e. Estimate the radii of three atoms not given in the data. Explain how you can estimate these values from the data given.

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9. The data below gives the ionic radii for several isoelectronic ions.

Ion Shell Core Charge Radius (pm)

S2– n = 3 +6 184

Cl– n = 3 +7 181

K+ n = 3 +9 133

Ca2+ n = 3 +10 104

a. Chemical species that have identical numbers of electrons are isoelectronic. How

many electrons do each of the ions have (total)?

b. What is the relationship between core charge and radius?

c. Predict which is larger: the O2– ion or the F– ion. Explain your reasoning.

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10. The data below gives the atomic radii of various atoms and ions.

Atom or Ion Shell Core Charge Valence e– Radius (pm)

F n = 2 +7 7 64

F– n = 2 +7 8 133

O n = 2 +6 6 66

O2– n = 2 +6 8 140

a. F and F– have identical core charges and the valence shell is n = 2 in both cases.

Explain why the radius of F– is larger than the radius of F.

b. Predict a value for the radius of N3–.

c. What are the three characteristics of an atom (or ion) that must be considered in determining relative radius?

Reflection Question: 1. Definitions for two more periodic trends are given below. Define these terms and

predict their general trend down a column and across a row (left to right) on the periodic table. Support your answer using the concepts covered in this section.

Electron Affinity: a measure of the likelihood of a neutral atom gaining an electron. It is the change in energy when an electron is added to a neutral atom to form an anion. Electronegativity: a measure of how well an atom attracts the electrons shared in a bond with another atom. For example, in an F‒F bond, the electrons are shared equally. In an H‒F bond, the electrons are shared unequally, closer to the F atom.

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1.8 Valence Electrons and Ions

Outcome: I can explain the relationship between trends in the reactivity of elements and periodicity.

Valence Electrons and Ionic Charge

1. For each statement, indicate if you agree or disagree and provide evidence to support

your position. a. Elements in group 1 (Li, Na, K) each have an s1 valence electron.

b. Elements in group 13 (B, Al, Ga) each have 2 valence electrons.

c. When Mg forms an ion, the charge will be −2.

d. If sodium reacts with any element in group 17 (F, Cl, Br), it will form a compound with a 1:1 ratio of each element.

e. If sodium reacts with any element in group 16 (O, S, Se), it will form a 1:1 ratio of each element.

f. If calcium and oxygen react, the compound formed is Ca2O.

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Alkali Metal Reactions

2. Observe the reaction (https://bit.ly/3mZtS1D) that occurs when alkali metals are exposed to air and when they are added to water. Notice that “this gradually becomes more terrifying as we go down the group”. Compare the metals to their position on the periodic table.

3. Metals react by losing electrons to form cations. Use this information, along with periodic knowledge, to explain the trend in reactivity of alkali metals.

Reflection Question: 1. Group 2 metals (alkaline earth metals) are known to form chloride salts in the form of

XCl2 (MgCl2, CaCl2, etc.). How does this finding support our understanding of electronic structure?

https://www.youtube.com/watch?v=uixxJtJPVXk

unit 1: atomic structure and properties

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