Summary

• In this unit, you will explore the fascinating world of the periodic table and uncover the patterns and trends that govern the behavior and properties of elements. Understanding these periodic trends is essential for predicting how elements will interact in chemical reactions and forming the basis for much of modern chemistry.

Key Concepts

1. Periodic Table:

   • The periodic table is a systematic arrangement of chemical elements based on their atomic number (the number of protons in the nucleus). It is divided into periods (horizontal rows) and groups (vertical columns).

2. Atomic Number:

   • The atomic number uniquely identifies each element and determines its position in the periodic table. It increases as you move from left to right across a period and from top to bottom down a group.

3. Atomic Radius:

   • Atomic radius is the size of an atom, typically defined as half the distance between the nuclei of two bonded atoms of the same element. It generally decreases across a period and increases down a group.

4. Ionization Energy:

   • Ionization energy is the energy required to remove an electron from an atom or ion. It tends to increase across a period and decrease down a group.

5. Electronegativity:

   • Electronegativity measures an atom's ability to attract electrons when participating in chemical bonds. It also increases across a period and decreases down a group.

6. Effective Nuclear Charge:

   • The effective nuclear charge is the net positive charge experienced by an electron in an atom, accounting for shielding by inner electrons. It affects many periodic trends.

Periodic Trends

• Atomic Radius Trend:

  • Atomic radius generally decreases from left to right across a period due to increased effective nuclear charge, which pulls electrons closer to the nucleus. It increases down a group as new energy levels (shells) are added.

• Ionization Energy Trend:

  • Ionization energy increases across a period(always left to right) because it becomes harder to remove electrons from smaller atoms that have more attractions

  • It decreases down a group because electrons are farther from the nucleus.

• Electronegativity Trend:

  • Electronegativity increases from left to right across a period, as atoms try to attract electrons more strongly.

  • It decreases down a group due to increased atomic size.

Practical Applications

• Understanding periodic trends helps predict chemical reactivity. Elements in the same group tend to have similar properties and form similar compounds.

• Ionization energy and electronegativity are crucial for understanding the behavior of elements in chemical bonding, such as the formation of ionic and covalent compounds.

• Atomic radius influences the size and geometry of molecules and affects properties like boiling and melting points.

Conclusion

Periodic trends are a cornerstone of chemistry, providing a systematic framework for understanding the properties and behaviors of elements. Mastery of these trends is essential for success in more advanced chemistry studies and applications in fields such as materials science, chemical engineering, and environmental science.

periodic law

• If the elements are organized according to atomic number, their physical and chemical properties will repeat.

  • Elements of the same group will have the same number of valence electrons and similar physical and chemical properties.

Simple concept overview

• The Periodic Law is like a secret code that helps chemists understand the behavior of elements.

• Imagine elements as different kinds of building blocks. When we line them up in a special order, we start to see that some properties,

  • like how big or small they are, how much they want electrons, or how they react with other elements, repeat in a predictable way.

  • It's like a musical rhythm in the world of atoms and helps us make sense of chemistry.

Complex

• When you move down a group (vertical column), you're adding more electron shells (energy levels), which affects atomic size, but other properties like reactivity with certain elements tend to repeat as well.

• properties of elements repeat in a predictable manner when arranged in the order of

  • increasing atomic number

• periodic law looks at the organization of elements based on their electron configurations(1s^1, 2s^1, etc) and the arrangement of electrons in atomic orbitals

• As you go across the period(left to right), the atomic number increases, which leads to a greater positive atomic charge

• this is because of trends such as atomic size, increasing electronegativity, and higher ionization energy

• These laws allow chemists to se elements behavior and properties

Atomic radius

• Size of the atom

• distance of two identical nuclei/ 2

• Average distance in an atom between the nucleus and the outermost electron(of the ATOM not the ion)

• Noble gasses are always the smallest

• Trends

  • Higher energy level is further from the nucleus

  • More charge on the nucleus pulls electrons in.

• Trends down a group

  • The attractive force of the nucleus of an atom on valence electrons.

  • As you move down a group, number of occupied Energy levels increases . The valence electrons are thereby further from the nucleus. Which creates a weakened effect on the core charge, which means it weakens the force of attraction.

• Trends across a period

  • The energy levels remain the same across a period.

  • Nuclear charge is based on the number of protons in the nucleus. Protons in the nucleus increasing, causing the charge to increase.This increases the attraction, which decrease the size of the radius.

Simple

• Atomic radius is like the size of an atom. Imagine an atom as a tiny planet with a nucleus at the center and electrons orbiting around it.

• The atomic radius tells us how far those outermost electrons are from the nucleus.

• Bigger atoms have electrons that orbit farther away and smaller atoms keep their electrons closer.

• it affects how atoms interact in chemical reactions.

Complex

• Atomic radius is a measure of the size of an atom, specifically the distance from the nucleus to the outermost electron shell.

• number of electron shells, the effective nuclear charge (the net positive charge experienced by the outermost electrons, taking into account shielding by inner electrons), and electron repulsion all influence atomic radius

• as you move down a group (vertical column) in the periodic table, atomic radius increases because additional electron shells are added, increasing the distance between the nucleus and the outermost electrons.

• as you move across a period (horizontal row) from left to right, atomic radius tends to decrease due to greater effective nuclear charge pulling the electrons closer to the nucleus.

• predicting how atoms bond and interact in chemical reactions, as it influences factors like bond length and reactivity.

• electrons are arranged in energy levels(shells)

Coulumbs law

• Columbic force of attraction equals charge on particle 1 and charge on particle 2 divided by radius (Distance between the particles)

• the attraction of [protons and electrons

• the more electrons you have the larger the charge

• Laymens terms; Coulomb's Law is like a rule that describes how electric charges behave. Imagine you have two charged objects, like magnets with different ends. Coulomb's Law tells you that the force between these charges depends on two things: how big the charges are and how far apart they are. If the charges are bigger, they pull harder or push harder. If they're closer, they also pull or push harder. It's like saying the more you charge an object (positive or negative), and the closer you bring it to another charged object, the stronger the attraction or repulsion between them.

• Complex terms; Coulomb's Law is a fundamental principle in physics that quantifies the force of attraction or repulsion between two point charges. It states that the magnitude of the electrostatic force (F) between two charges (q₁ and q₂) is directly proportional to the product of their charges and inversely proportional to the square of the distance (r) separating them.

Ionic radius

• Size of the ion

• Average distance between nucleus and outermost electron of the Ion (has a different number of electrons)

• Cations (+, generally metals) have a smaller radius whereas Anions (- generally non-metals) have a larger radius.

• noble gasses are always the smallest in ionic radius

electronegativity

• adding an electron

• opposite of ionization energy

Simple

• the greediness factor for atoms in a molecule

• tells us how much an atom in a molecule wants to hog electrons

• high electronegativity = wants the electrons more

• low electronegativity = wants the electrons less(and shares them more easily

Complex

• the ability to attract electrons towards itself(the atom)

• helps us undertsand how electrons are shared or pulled between atoms

PES graphs

• Atoms= nesting dolls

• each doll represents and electron shell

• Stands for “potential energy surface”

• what we can learn from a PES graph

  • energy levels - height of the top points of the graph tells you how much pointial energy it has

  • higher hills mean higher energy and vice versa

  • stability - the lowest points on the graph represent the most stable configurations(this is the area the atom “wants” to be in b/c it has the least amount of potential energy

  • reations - How atoms rearrange themselves throughout a chemical reaction

    • the reactants will be at the start and products will be at the back

    • the reaction occurs when the atom is moved from higher energy states to lower energy states

  • activation energy - The energy difference between reactants starting point and the highest point

• How to identify an element

  • look at the peaks and write out electron configuration

Ionization energy

• amount of energy it takes to remove an electron

Simple

• Think of an atom like a group of balloons. Each electron in the atom is like a balloon tied to a string. These balloons (electrons) are floating around the atom. Now, if you want to take one of those balloons (electrons) away from the atom, you need to tug on the string with enough force to break it and make the balloon float away on its own.

• ionization energy is the strength of those strings

• how tightly an atom holds to its electrons

Complex

• A measure of the strength of attraction between the negatively charged electrons and the positively charged nucleus

• electrons in the inner most shells are held on to most tightly

• outer shells are easier to remove

• the ionization energies get increasingly higher as you continue to remove electrons

• the atom can “ionize”(hence ionization energy)

• How to identify an element based off a chart

  • look at all the ionization energies

  • find the biggest jump

  • wherever that jump is is the number of valence electrons

    • for example, if the biggest jump was between the second ad third ionization energy, it would have 2 valence electrons

Vocabulary

1. Periodic Table: A tabular arrangement of chemical elements, organized by their atomic number and grouped into periods (horizontal rows) and groups or families (vertical columns).

2. Atomic Number: The number of protons in the nucleus of an atom, which uniquely identifies each element and determines its position in the periodic table.

3. Period: A horizontal row in the periodic table, representing elements with the same number of electron shells (energy levels).

4. Group/Family: A vertical column in the periodic table, comprising elements with similar chemical properties due to having the same number of valence electrons.

5. Atomic Radius: The size of an atom, typically defined as half the distance between the nuclei of two bonded atoms of the same element.

6. Ionization Energy: The energy required to remove an electron from an atom or ion, indicating how strongly an atom holds its electrons.

7. Electronegativity: A measure of an atom's ability to attract electrons when participating in chemical bonds; higher electronegativity means greater attraction.

8. Shielding Effect: The reduction in the effective nuclear charge (positive charge experienced by outer electrons) due to inner electrons that repel and shield the outer electrons from the nucleus.

9. Valence Electrons: Electrons in the outermost energy level of an atom, which are primarily responsible for the chemical behavior of an element.

10. Noble Gases: Elements in Group 18 (VIII A) of the periodic table, known for their stable electron configurations and low reactivity.

11. Alkali Metals: Elements in Group 1 (I A) of the periodic table, highly reactive metals that have one valence electron.

12. Alkaline Earth Metals: Elements in Group 2 (II A) of the periodic table, reactive metals with two valence electrons.

13. Transition Metals: Elements found in the central block of the periodic table (Groups 3-12), known for their variable oxidation states and diverse properties.

14. Halogens: Elements in Group 17 (VII A) of the periodic table, highly reactive nonmetals that have seven valence electrons.

15. Metalloids: Elements with properties intermediate between metals and nonmetals, often found along the staircase on the periodic table.

16. Effective Nuclear Charge: The net positive charge experienced by an electron in an atom, accounting for shielding by inner electrons.

17. Periodic Trend: A predictable pattern or change in properties of elements across the periodic table, such as atomic size, ionization energy, and electronegativity.

These terms are fundamental to understanding how elements are organized and how their properties vary systematically in the periodic table, which is crucial for comprehending chemical behavior and reactions.

Multiple-Choice Questions:

1. Which of the following elements has the highest electronegativity? a. Sodium b. Chlorine c. Potassium d. Calcium

2. As you move across a period from left to right in the periodic table, what generally happens to the atomic radius? a. It increases. b. It remains constant. c. It decreases. d. It fluctuates unpredictably.

3. Which of the following elements has the highest ionization energy? a. Lithium b. Carbon c. Oxygen d. Neon

4. What is the trend in electron affinity as you move down a group in the periodic table? a. It increases. b. It decreases. c. It remains constant. d. It depends on the specific group.

5. Which element has the highest first ionization energy? a. Sodium b. Magnesium c. Aluminum d. Silicon

6. Among the noble gases, which has the highest boiling point? a. Helium b. Neon c. Argon d. Xenon

7. What happens to the metallic character of elements as you move from right to left across a period? a. It increases. b. It decreases. c. It remains constant. d. It varies randomly.

8. Which of the following elements is the least reactive with other elements? a. Fluorine b. Chlorine c. Bromine d. Iodine

9. Which element is found in the same group as oxygen? a. Nitrogen b. Sulfur c. Carbon d. Neon

10. What is the most abundant element in Earth's crust? a. Oxygen b. Silicon c. Aluminum d. Iron

11. Which element is known for its extreme reactivity with water and air? a. Sodium b. Phosphorus c. Potassium d. Calcium

12. Which of the following elements is a metalloid? a. Boron b. Sodium c. Iodine d. Argon

13. What is the trend in atomic radius as you move down a group in the periodic table? a. It increases. b. It decreases. c. It remains constant. d. It fluctuates unpredictably.

14. Which element has the highest electron affinity? a. Fluorine b. Chlorine c. Bromine d. Iodine

15. What happens to the atomic radius of an element as you move from top to bottom within a group? a. It increases. b. It decreases. c. It remains constant. d. It fluctuates unpredictably.

16. Which element has the highest melting point? a. Carbon b. Boron c. Tungsten d. Mercury

17. Which element is considered the least electronegative? a. Francium b. Lithium c. Fluorine d. Chlorine

18. What is the trend in ionization energy as you move from left to right across a period? a. It increases. b. It decreases. c. It remains constant. d. It fluctuates unpredictably.

19. Which element is a transition metal in Period 4? a. Titanium b. Potassium c. Calcium d. Chlorine

20. Among the alkali metals, which one has the highest atomic number? a. Lithium b. Sodium c. Potassium d. Rubidium

21. Which element is known for its gaseous state at room temperature? a. Oxygen b. Nitrogen c. Hydrogen d. Carbon

22. What is the trend in electron affinity as you move from left to right across a period? a. It increases. b. It decreases. c. It remains constant. d. It fluctuates unpredictably.

23. Which element is a noble gas in Period 3? a. Neon b. Argon c. Krypton d. Xenon

24. What is the trend in metallic character as you move down a group in the periodic table? a. It increases. b. It decreases. c. It remains constant. d. It fluctuates unpredictably.

25. Which element is a halogen in Period 2? a. Fluorine b. Chlorine c. Bromine d. Iodine

26. What is the trend in atomic radius as you move from right to left across a period? a. It increases. b. It decreases. c. It remains constant. d. It fluctuates unpredictably.

27. Which element has the highest first ionization energy? a. Sodium b. Magnesium c. Aluminum d. Silicon

28. Which element is a noble gas in Period 4? a. Neon b. Argon c. Krypton d. Xenon

29. Among the alkaline earth metals, which one has the lowest atomic number? a. Beryllium b. Magnesium c. Calcium d. Strontium

30. Which element is known for its shiny appearance, malleability, and good conductor of heat and electricity? a. Oxygen b. Sulfur c. Copper d. Neon

Free-Response Questions (FRQs):

31. Explain the trend in atomic radius as you move from left to right across a period in the periodic table.

32. Describe the variations in metallic character as you move down a group in the periodic table.

33. Discuss the factors that influence the ionization energy of an element.

Answer Key:

1. b

2. c

3. d

4. b

5. c

6. d

7. b

8. d

9. b

10. a

11. c

12. a

13. a

14. a

15. a

16. c

17. a

18. a

19. a

20. d

21. c

22. a

23. b

24. a

25. a

26. b

27. d

28. d

29. a

30. c

FRQ Answers:

31. Atomic radius generally decreases as you move from left to right across a period. This is because the increasing number of protons in the nucleus exerts a stronger positive charge on the electrons in the same energy level, pulling them closer to the nucleus.

32. Metallic character increases as you move down a group in the periodic table. This

1. Which element in Group 17 (Group VIIA) has the highest electronegativity? a) Fluorine b) Chlorine c) Bromine d) Iodine

2. What trend explains why atomic radii generally decrease across a period in the periodic table? a) Shielding effect b) Effective nuclear charge c) Electron affinity d) Ionization energy

3. Among the alkali metals, which element has the highest ionization energy? a) Lithium b) Sodium c) Potassium d) Rubidium

4. Which of the following elements is a noble gas? a) Neon b) Phosphorus c) Sulfur d) Chlorine

5. What happens to the atomic radius as you move down a group (family) in the periodic table? a) It decreases b) It remains constant c) It increases

6. Which element is the least metallic in Group 14 (Group IVA)? a) Carbon b) Silicon c) Germanium d) Lead

7. Which property of elements increases as you move from left to right across a period? a) Electronegativity b) Atomic radius c) Ionization energy d) Electron affinity

8. Among the halogens, which element has the largest atomic radius? a) Fluorine b) Chlorine c) Bromine d) Iodine

9. What is the primary reason behind the increase in ionization energy as you move from left to right across a period? a) Decreased effective nuclear charge b) Increased shielding effect c) Greater number of protons d) Larger atomic radius

10. Which element in Group 2 (Group IIA) has the highest melting point? a) Beryllium b) Magnesium c) Calcium d) Barium

Free-Response Questions:

11. Explain the trends in atomic radius across a period and down a group in the periodic table. Provide reasons for these trends.

12. Discuss the variations in ionization energy as you move across a period. Explain the underlying factors responsible for these changes.

13. Compare and contrast the properties of alkali metals and noble gases based on their positions in the periodic table.

Answer Key:

1. a) Fluorine

2. b) Effective nuclear charge

3. d) Rubidium

4. a) Neon

5. c) It increases

6. a) Carbon

7. a) Electronegativity

8. d) Iodine

9. c) Greater number of protons

10. c) Calcium

Free-Response Answers:

11. (Answers may vary but should include discussion of effective nuclear charge, shielding, and electron configurations.)

12. (Answers should include explanations involving effective nuclear charge, shielding, and electron configurations.)

13. (Answers should highlight differences in reactivity, electron configurations, and physical properties between alkali metals and noble gases.)

Feel free to create more questions or expand on the explanations in the free-response section as needed for your specific study guide. This should provide a challenging foundation for your honors chemistry study guide on periodic table trends.

Redo worksheets but you mustn’t look at notes, they are already here ;)

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