The Periodic Table Chapter 4

4.1 Mendeleev’s Periodic Table (1869) Reading Check Why did Mendelev place certain elements together in the same group? Arranged elements with similar properties into groups.

4.1 Mendeleev’s Periodic Table (1869) Why did Mendeleev leave gaps in his periodic table? ANS = He predicted that elements would later be discovered to occupy these gaps

4.1 Periodic Table Why does the PT have the shape it does?

4.1 Periodic Table Elements are organized in order of increasing atomic number Atomic mass follows the same order (in most cases)

4.1 Periodic Table Mendeleev did not know why certain elements have similar properties because he did not know about atomic theory. But we do now! Why do properties repeat themselves? The properties of an element are determined by it’s valence electrons.

3.3 Valence Electron Anelectronthatisfoundontheoutermostshellofanatomandthatdeterminestheatom’schemicalproperties.

3.3 Valence Electron An electron that is found on the outermost shell of an atom and that determines the atom’s chemical properties. The electrons that are on the inside are called core electrons Which element is this? 

Think about it….. How many valence electrons are in there in a sulfur atom? (first, write the electron configuration)

Think about it….. 1s2 2s2 2p6 3s2 3p4 2-8-6 How many valence electrons are in there in a sulfur atom? 6 valence electrons The rest are core electrons. How many core electrons? 10

4.1 Valence electrons How many valence electrons are there in a fluorine atom? (first, write the electron configuration for fluorine)

4.1 Valence electrons How many valence electrons are there in a fluorine atom? 1s22s22p5 [He]2s22p5 2-7

4.1 Periodic Table A period (group/period) is a horizontal row in the periodic table  numbered 1-7 Corresponds to the energy level of the valence electrons

4.1 Periodic Table A group is a vertical column in the periodic table ↓ Groups 1-18 See fig 3 p 119 s-block, p-block, d-block, f-block

4.3 Trends in the periodic table Standard Use the periodic table as a model to predict the relative properties of elements based on the patterns of electrons in the outer energy level of atoms.

4.3 Trends in the periodic table Learning Target Students learn about periodic trends in atomic radius by plotting/interpreting a graph

Think about it…. Atoms are small……but are they all the same size? Is a carbon atom the same size as an iron atom?

4.3 Trends in the PT Worksheet: Family Characteristics Part 1: Atomic Radius How is atomic radius defined? See fig 19 p. 135

4.3 Atomic Radius fig 21 p. 136 Inference: What is this diagram telling us?

4.3 Trend in Atomic Radius Draw in the lines for groups 1,2,13,14, 17,18 on your worksheet. Inference: What is your graph telling you? What is the trend in atomic radius as you read down a group in the PT? What is the trend in atomic radius as you read across a period in the PT?

4.3 Trends in the PT How does atomic radius change as you read down a group in the PT? ANS = Atomic radius increases (increases/decreases).

4.3 Atomic radius Notes: Explain why (1 sentence)

4.3 Trends in the PT How does atomic radius change as you read across a period in the PT? ANS = Atomic radius decreases (increases/decreases).

4.3 Trends in the PT COUNTERINTUITIVE  Factors: Effective nuclear charge Shielding Read p. 136 ‘Atomic Radius Decreases as You Move Across a Group’

4.3 Electron Shielding Factors: Effective nuclear charge Shielding Calculate effective nuclear charge for Na and Ar Factors: Effective nuclear charge Shielding

4.3 Atomic Radius Summary fig 20 p. 135 Inference: where are the largest elements found on the PT? The smallest?

4.3 Atomic Radius Summary fig 21 p. 136 Inference: what is this diagram telling us? How does this diagram differ from the one we plotted?

Check For Understanding: Atomic Radius Which of the following elements has the smallest atomic radius? Magnesium Calcium Strontium Barium 2. Which of these elements has the largest atomic radius? Gallium Germanium Selenium Bromine

4.3 Trends in the periodic table Standard Use the periodic table as a model to predict the relative properties of elements based on the patterns of electrons in the outer energy level of atoms. Learning Target Students learn about periodic trends in ionization energy by plotting/interpreting a graph

4.3 Ions Notes: Whenatomshaveequalnumbersof protonsandelectronstheyareelect ricallyneutral

4.3 Ions When atoms have equal numbers of protons and electrons they are electrically neutral. Butwhenenoughenergyisadded,t heattractiveforcecanbeovercome.

4.3 Ions But when enough energy is added, the attractive force can be overcome. Whenthishappens,anelectronisre movedfromtheatom,anditbecome sapositivelychargedion

4.3 Ions When this happens, an electron is removed from the atom, and it becomes a positively charged ion. (p. 133) Positively charged ion = ‘cation’

IONS

Na atom  Na+ ion + electron 4.3 Ionization Energy Ionization Energy = The energy that is needed to remove an electron from an atom. Na atom  Na+ ion + electron

Ionization Energy Unit = (kJ/mole) ‘First Ionization energy’ = energy needed to remove the first electron from an atom

4.3 Ionization energy If ionization energy is high, it is difficult (difficult/easy) to remove an electron from the atom.

4.3 Ionization energy If ionization energy is low, it is easy (difficult/easy) to remove an electron from the atom.

4.3 Trend in ionization energy Read Family Characteristics worksheet Part 2 Draw in the lines for groups 1,13,15,17,18 on your worksheet. Inference: What is your graph telling you?

4.3 Ionization Energy Ionization Energy decreases (increases or decreases) as you read down a group in the PT. Ionization Energy increases (increases/decreases) as you read across a period in the PT

Think about it….. In the Brainiac alkali metals video clip, which group 1 element was the most reactive when put into water? ANS = Those at the bottom Rb & Cs

4.3 Ionization Energy ‘Ionization Energy decreases as you read down a group in the PT.’ Which elements in group 1 have the lowest ionization energy? ANS= Those at the bottom Rb & Cs

4.3 Electron Shielding Notes: Explain why (1 sentence) Read p. 133 ‘Ionization Energy Decreases As You Move Down a Group’

4.3 Effective Nuclear Charge Why is the ionization energy for chlorine greater than that for magnesium? Read Ionization Energy Increases as you Move Across a Period p. 134

4.3 Ionization Energy See fig 17 p. 134 Inference: which elements have the highest ionization energy? The lowest?

Periodic Trends What is the relationship between atomic radius and ionization energy? An atom with a small atomic radius has a high (high/low) ionization energy. e.g. Helium

Periodic Trends What is the relationship between atomic radius and ionization energy? An atom with a large atomic radius has a low (high/low) ionization energy. e.g. Cesium

Check For Understanding: Ionization Energy Which of the following elements has the lowest ionization energy? Magnesium Calcium Strontium Barium 2. Which of these elements has the highest ionization energy? Gallium Germanium Selenium Bromine

Electronegativity Definition: Ameasureoftheabilityofanatominachemicalcompoundtoattractelectronstoitself. A measure of the ability of an atom in a chemical compound to attract electrons to itself. (p. 137) Facts: Read/Notes ‘Electronegativity’ p. 137 (first 3 paragraphs)

think about it…….. Periodic Trends What is this diagram telling us? How does ‘electronegativity’ vary as we move around the PT? Scale: 0-4

Electronegativity See p. 137 fig 22 Where are the most electronegative elements found in the PT? Where are the least electronegative elements found in the PT?

5.1 Electronegativity Fluorine has the strongest demand for electrons Cesium has the lowest demand for electrons

Periodic Trends What is the relationship between ionization energy and electronegativity? An atom with a high ionization energy has a high (high/low) electronegativity e.g. Fluorine Directly proportional

5.1 Electronegativity Electronegativity difference between two atoms tells us what kind of bond will form between them.

5.1 Electronegativity Calculate the EN difference between sodium and chlorine = Ionic Bond

5.1 Electronegativity Electronegativity difference between two atoms tells us what kind of bond will form between them.

5.1 Electronegativity Calculate the EN difference between sulfur and oxygen = covalent bond

Other Periodic Trends p. 139 Electron Affinity Melting and Boiling Points