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Chapter 6 Lesson 3 “Periodic Trends” The Big Idea…  Nuclear Charge – the effect protons of an atom have on its size and shape.

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Presentation on theme: "Chapter 6 Lesson 3 “Periodic Trends” The Big Idea…  Nuclear Charge – the effect protons of an atom have on its size and shape."— Presentation transcript:

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2 Chapter 6 Lesson 3 “Periodic Trends”

3 The Big Idea…  Nuclear Charge – the effect protons of an atom have on its size and shape.

4 I. Trends in Atomic Size A. Atomic Radius 1. Def – one half of the distance between the nuclei of two like atoms in a diatomic molecule. -radius = from center to outer edge 2. Measured in picometers (pm) -1 x m 3. Measured in angstroms (Å) -1 x m

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6 3. The distance between the nuclei of two covalently bonded flourine atoms is 128 pm. What is the atomic radius of 1 fluorine atom? 4. The distance between the nuclei of two covalently bonded nitrogen atoms is 142 pm. What is the atomic radius of 1 nitrogen atom in nanometers? 5. The distance between the nuclei of two covalently bonded iodine atoms is 276 pm. What is the diameter of 1 iodine atom in pm? 64 pm.0710 nm 276 pm

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9 B. Group/Family Trends 1. Size increases as you move down a column. -Why? electrons are added to energy levels farther away from the nucleus. -How can electrons get into energy levels away from the nucleus? electrons are shielded from the positive nuclear charge by 1s electrons…2s electrons… 2p electrons… and so forth. 2. Shielding demo

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11 In Class Assignment 1. Place the following atoms in order from smallest to largest in terms of their atomic radius. Br, I, Be, He, Rn 2. Why is a sodium atom smaller than a potassium atom? He < Be < Br < I < Rn Because potassium has more electrons and these electrons fill energy levels that are farther away from the nucleus of a sodium atom.

12 C. Periodic Trends 1. Atomic radii decreases as you move from left to right. Why? electrons are added to the same principle energy level. Nuclear charge pulls electrons in closer to the nucleus.

13 Group #1 Group #2

14 Period #2Period #3

15 In Class Assignment 1. How would you describe the atomic radius of a period 2 alkaline earth metal compared to a period 4 alkaline earth metal? 2. How would you describe the atomic radius of a period 3 alkali metal and a period 3 halogen? The atomic radius of the period 2 a.e.m. would be smaller than the period 4 a.e.m. The atomic radius of a period 3 a.m. would be larger than a period 3 halogen due to nuclear charge.

16 II. Trends in Ionization Energy A. Def – the amount of energy required to overcome the attraction of the nuclear charge and remove an electron from an atom. Energy Na e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- 1+

17 1. 1 st Ionization Energy -removing the 1 st electron 2. 2 nd Ionization Energy -removing an electron from a 1 + ion rd …4 th …5 th … and so on **Note table 6.1

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20 B. Group Trends 1. Ionization energy decreases as you move down a group. Why? The farther electrons are from the nuclear charge, the easier they are removed. C. Period Trends 1. Ionization energy increases as you move from left to right. Why? The nuclear charge increases, but the electron shield does not. **Remember noble gases do not want to give up electrons**

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22 III. Trends in Ionic Size A. Ion Review 1. metals  low ionization energy 2. non-metals  high ionization energy **How does losing/gaining an electron affect the size of an ion???** B. Cations and Anions 1. Cations = always smaller than their neutral atoms Why? loss of an electron causes the nuclear charge to increase. Thus the remaining electrons are pulled in farther.

23 e-e- e-e- e-e-

24 2. Anions = always larger than their neutral atom. Why? The increase of another electron causes the nuclear charge to decrease. Thus the size of the ion increases.

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26 C. Group Trends 1. Ionic size increases as you move down a group on the periodic table. Increases Cations decrease Anions decrease D. Periodic Trends 1. Ionic size decreases from left to right for both cation and anions

27 1. How does the ionic radius of sodium compare to that of cesium? 2. How does the ionic radius of boron compare to that of fluorine? Cesium is larger Fluorine is larger

28 IV. Electronegatvity A. Def – the tendency for an atom to attract electrons from another atom. 1. similar to magnetism 2. Electronegativity is a relative value based on ionization energy. 3. Noble Gases are not included…Why??? -they don’t form compounds B. Electronegativity Trends 1. E-negativity decreases as you move down a column. 2. E-negativity increases as you move across a period. -Fluorine = most electronegative element -Cesium = least electronegative element

29 3. Metals = low e-negativity Non-metals = high e-negativity E-negativity = tug of war Fluorine or Cesium? Calcium or Sulfur? Oxygen or Magnesium? Oxygen or Nitrogen? 4. E-negativity values help predict types of bonds. Fluorine Sulfur Oxygen

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