Periodic Trends 6.3

1. Periodic Trends6.3

2. Atomic Radius • The distance between an atom’s nucleus and its valence electrons. • How closely an atom lies to a neighboring atom • Size of the atom varies from substance to substance

3. Trends within Periods • Atomic Radii DECREASE as you move left to right across a period • Because of Increased nuclear charge(total charge of all protons in the nucleus) • Increased nuclear charge pulls the outermost electrons closer to the nucleus & decreases the atomic radius.

4. Trends within Groups • Atomic Radii INCREASE as you move down a group. • Nuclear charge increases and electrons are added to higher energy levels. • Moving down a group: • The Outermost orbital increases in size • Increasing principal energy level

5. Examples • Which element has the smallest atomic radius? Largest atomic radius? Iodine (I) Bromine (Br) Fluorine (F) Chlorine (Cl) • Fluorine • Iodine

6. Ionic Radius • Ion- an atom that gains or loses electrons • When atoms lose electrons they form positive ions and become smaller • The electron lost will always be a valence electron • Loss of valence electrons may leave an empty outer orbital which results in a smaller radius • Repulsion between fewer electrons decreases allowing them to be pulled closer to the nucleus

7. When atoms GAIN electrons they form negative ions and they always become larger. • Increase of an electron to an atom increases the repulsion between the valence electrons forcing them to move farther apart. • This equals larger radius

8. Ionic Radius within Periods • Size of the positive ions gradually decreases • Beginning in group 5A or 6A the size of the much larger negative ions also gradually decreases. • Ionic Radii DECREASE across periods

9. Ionic Radius within Groups • As you move down a group an ion’s outer electrons are in higher principal energy levels resulting in a gradual increase in ionic size • Ionic Radii increase as you move down a group

10. Ionization Energy • The energy required to remove an electron from an atom • “How strongly an atom’s nucleus holds on to its valence electrons” • High IE indicates atom has a strong hold on its electrons • Low IE indicates an atom loses its outer electron easily

11. Ionization Energy • Energy required to remove the 1st electron is the first ionization energy. • Energy required to remove the 2nd electron is the second ionization energy.

12. Ionization Energy within Periods • INCREASE as you move from left to right across a period • The increased nuclear charge of each successive element produces an increased hold on the valence electrons.

13. Ionization Energy within Groups • DECREASE as you move down a group • Occurs because atomic size increases as you move down a group • With the valence electrons farther from the nucleus less energy is required to remove them

14. Octet Rule • Atoms tend to gain, lose, or share electrons in order to acquire a full set of eight valence electrons. • Elements on the right side of the periodic table tend to gain electrons in order to acquire the 8 valence electrons. (Form negative ions) • Elements on the left side of the periodic table tend to lose electrons and form positive ions.

15. Electronegativity of an Element • Indicates the relative ability of its atoms to attract electrons in a chemical bond • Noble Gases are not assigned values • Fluorine is the most electronegative • Fr & Cs are the least • In a chemical bond the atom with the greater electronegativity more strongly attracts the electrons

16. Electronegativity Trends within Periods & Groups • INCREASES as you move across a period • DECREASES as you move down a group • The lowest electronegativities are found at the lower left side • Highest are found at the upper right side

17. Example Problems • Which element has the highest electronegativity? Lowest? • N- Nitrogen • P- Phosphorus • As-Arsenic • Sb-Antimony • Bi- bismuth • N=highest • Bi= Lowest

18. Homework Problems • Pg. 175 • #56, 57, 59, 62, & 63