Shapes of the orbitals

Transcription

1 Electrons Review and Periodic Table Trends Unit 7

2 Electrons Shapes of the orbitals

3 Electron Configuration Electrons spin in opposite direction

4 Background Electrons can jump between shells (Bohr s model supported) The electrons can be pushed so far that they escape the attraction of the nucleus Losing an electron is called IONIZATION Remember an ion is an atom that has either a net positive or net negative charge Q: What would the charge be on an atom that lost an electron? Gained two electrons?

5 IONIZATION ENERGY o Defn: Ionization energy is the energy required to remove one outer electron from an atom o When an electron is taken away, what kind of ion results? o A positively charged ion (cation) o Review: Oxidation numbers Metals form (+) or (-) ions? + Nonmetals form (+) or (-) ions? - FIRST IONIZATION ENERGY - the energy required to pull off the first valence electron. SECOND IONIZATION ENERGY - the energy required to pull off the second valence electron.

6 Ionization Energy cont. Low IE = not a lot of energy required = cations formed easily High IE = takes a lot of energy = cations NOT formed; anions What has a higher ionization energy a metal or nonmetal? Nonmetal Why? Tend to form anions

7 IONIZATION ENERGY a) Moving left to right across the periodic table IONIZATION ENERGY INCREASES (harder to pull off an electron) WHY? More protons are in the nucleus therefore the valence electrons are strongly attracted to the nucleus which increases the energy required to remove them. b) Moving down a group, IONIZATION ENERGY DECREASES (easier to pull off an electron) WHY? Shielding effect Shielding - core e- block the attraction between the nucleus and the valence e- c) The second ionization energy is greater, third is even greater WHY? Electrons that remain move closer to the nucleus.

8 The ionization energy is the amount of energy needed to strip an electron off of an atom, ion, or molecule. The illustration shows that the metals like lithium, Li, and cesium, Cs, have relatively low ionization energies. This means it takes relatively small amounts of energy to remove electrons from these atoms. Metals tend to lose electrons and form positive ions. The nonmetals like neon, Ne, fluorine, F, and oxygen, O, have relatively high ionization energies. This indicates that the nonmetals have strong attractions for their valence electrons. The nonmetals hold on to their electrons. In fact nonmetals gain electrons to form negative ions.

9 Ionization energy (kj/mol) H He Li Be B C N O F Ne Na Mg Al Si P S Ar Cl K Ca Element

10 Factors Affecting Ionization Energy

11 ATOMIC RADIUS o The approximate distance from the nucleus of an atom to its valence electrons. a) Moving left to right across the periodic table ATOMIC RADIUS DECREASES WHY? More protons are being added to the nucleus, valence electrons are strongly attracted to the nucleus. Electrons are also being added, but in the same shell at about the same distance so there is not much of a shielding effect.

12 Atomic Radius Cont. H b) Moving down a group, ATOMIC RADIUS INCREASES Li Na WHY? Each atom has another energy level, so the atoms get bigger. Since electrons are being added to distant shells away from the nucleus the valence electrons are SHIELED by the inner shell electrons. K Rb

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14 Atomic Radius (pm) 250 K Li Be B Na C N O F Ne Mg Al Si P S Cl Ar Ca 0 H He Element

15 Review The size of an atom is largely determined by its electrons. The electrons are arranged in shells surrounding the nucleus of each atom. The top elements of every group have only one or two electron shells. Atoms of elements further down the table have more shells and are therefore larger in size. Moving across a period from left to right, the outermost electron shell fills up but no new shells are added. At the same time, the number of protons in the nucleus of each atom increases. Protons attract electrons. The greater the number of protons present, the stronger the attraction that holds the electrons closer to the nucleus, and the smaller the size of the shells.

16 ELECTRONEGATIVITY o Defn: How strongly the nucleus of an atom attracts the electrons of other atoms in a bond. o High electronegativity = wants to gain an electron = easy to become a negative ion o Given specific values and found on electronegativity table DO NOT NEED TO MEMORIZE! a) Moving from left to right on the periodic table ELECTRONEGATIVITY INCREASES. b) Moving down a group ELECTRONEGATIVITY DECREASES.

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19 Review of Electronegativity Atoms of different elements have different attractions for bonding electrons. Electronegativity is a measure of the attraction an atom involved in a bond has for the electrons of the bond. An atom with a high electronegativity will tend to attract bonded electrons towards it. An atom with a low electronegativity will have a very weak attraction for electrons. Electronegativity values can be useful in predicting which type of bonding is most likely between two elements

20 Electronegativity cont. What element has the highest EN? F Which would have greater EN, a metal or a nonmetal? Nonmetal Which is more electronegative, Cu or S? S Br or Ga? Br

21 Other Trends Melting Points Metals USUALLY decreases as you go down a group Non-metals USUALLY increases as you go down a group

22 F. Melting/Boiling Point Melting/Boiling Point Highest in the middle of a period

23 Other Trends Reactivity Defn: how likely or vigorously an atom is to react with other substances Usually determined by how easily electrons can be removed (ionization energy) and how badly they want to take other atom s electrons (electronegativity) since it is the transfer/interaction of electrons that is the basis of chemical reactions

24 Reactivity Cont. Metals Period: DECREASES from left to right Group: INCREASES down a group WHY? The farther left and down the periodic table, the easier it is for elections to be given or taken away, resulting in a higher reactivity Non-metals Period: INCREASES from left to right Group: DECREASES down a group WHY? The farther right and up the periodic table, the higher the electronegativity, resulting in a more vigorous exchange of electrons

25 Periodic Trend Summary Based upon an element s position on the periodic table, it is possible to make predictions regarding its behavior. All periodic trends can be understood in terms of three basic rules: 1.) Electrons are attracted to the protons in the nucleus of an atom a) The closer an electron is to the nucleus, the more strongly it is attracted. b) The more protons in a nucleus, the more strongly an electron is attracted.

26 Periodic Trend Summary Cont. 2.) Electrons are repelled by other electrons in an atom. So if other electrons are between a valance electron and the nucleus, the valence electrons will be less attracted to the nucleus. This is called the SHIELDING EFFECT. 3.) Completed orbits, and sub-orbits, are very stable. Atoms prefer to add or subtract valence electrons to create complete shells if possible.

27 Review: IE vs. EN A high IE tends to form a: (positive, negative) ion A high EN tends to form a: (positive, negative) ion Metals have a: (high, low) IE Nonmetals have a: (high, low) EN