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Honors Chemistry Chapter 5 Lesson 3

Honors Chemistry Chapter 5 Lesson 3. “Periodic Trends”. The Big Idea…. Nuclear Charge – the effect protons of an atom have on its size, shape and characteristics. NC +. Periodic Table. NC +. I. Trends in Atomic Size. A. Atomic Radius

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Honors Chemistry Chapter 5 Lesson 3

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  1. Honors ChemistryChapter 5 Lesson 3 “Periodic Trends”

  2. The Big Idea… • Nuclear Charge – the effect protons of an atom have on its size, shape and characteristics. NC + Periodic Table NC +

  3. 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 a. Picometers (10-12 m) b. angstroms (Å) (10-10 m) c. Nanometers (10-9)

  4. 3. The distance between the nuclei of two covalently bonded fluorine 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? 64.0 pm .0710 nm 276 pm

  5. 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. C. Periodic Trends 1. size decreases as you move left to right in a period -Why? electrons are shielded from the positive nuclear charge by 1s electrons…2s electrons… 2p electrons… and so forth. 2. Shielding demo

  6. 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.

  7. 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.

  8. 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 1+ e- e- e- e- e- e- e- e- e- Na Na e- e- e- e- e- e- e- e- e- e- e- e-

  9. 1. 1st Ionization Energy -removing the 1st electron 2. 2nd Ionization Energy -removing an electron from a 1+ ion. 3. 3rd…4th…5th… and so on **Note table in book…

  10. I.E. + I.E. - Periodic Table

  11. 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**

  12. III. Trends in Ionic Size **How does losing/gaining an electron affect the size of an ion???** A. Ionic Size 1. metals form cation low ionization energy a. Smaller than their neutral atom Why? loss of an electron causes the nuclear charge to increase. Thus the remaining electrons are pulled in farther.

  13. e- e- e-

  14. 2. non-metals form anions high ionization energy a. larger than their neutral atom. Why? The increase of another electron causes the nuclear charge to decrease. Thus the size of the ion increases.

  15. - Ionic Size Ionic Size +

  16. IV. Electron Affinity A. Def – the energy change that occurs when it gains an extra electron. 1. Attraction for extra electron(s) 2. Most have negative affinities pg. 182 3. loose energy when an electron is added B. Trends 1. Is there one? 2. Pg. 182 a) non-metals  more negative b) metals  less negative 3. Noble Gases  positive affinity *related to the number of e’s needed to octet

  17. V. Electronegativity A. Def – the tendency for an atom to attract electrons from another atom. 1. similar to magnetism 2. Scale on page 184 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 -Francium = least electronegative element

  18. 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 Oxygen

  19. Electronegativity Chart 4.0 strongest 0.7 weakest

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