9 Classification of the Elements Metals Nonmetals Metalloids
10 Occupy the left side of the periodic table Have luster, shiny MetalsOccupy the left side of the periodic tableHave luster, shinySolids at room temp except HgDuctile: ability to be drawn into wiresMalleable: ability to be hammered into sheetsExcellent conductors of heat and electricityTend to form positive ions
11 Occupy the right side of the Periodic Table NonMetalsOccupy the right side of the Periodic TableGenerally gases or brittle solidsDull-lookingBrittlePoor conductors of heat and electricityBromine is the only liquid at room tempTend to form negative ions
12 MetalloidsCharacteristics of metals and nonmetals
13 Classification of the Elements Families of elements share the same ending electron configurationtherefore they share similar chemical characteristicsValence Electrons: electrons in the highest principal energy levelDetermine Chemical reactivityElements in a group share the same number of valence electrons
15 Number of Valence Electrons Elements on the rightNonmetals4 or more valence electronstend to gain electronsbecome negative ionsElements on the leftMetals3 or less valence electronstend to lose valence electronsform positive ions
17 Families of elementsElements of the same family (group) share structural and chemical (behavioral) characteristicsAlkali MetalsAlkaline Earth MetalsTransition ElementsHalogensNobel Gases
18 Group 1: Alkali Metals Soft, highly reactive metals Usually stored under oil or kerosene to prevent their interaction with air and water
19 Properties of Alkali Metals React vigorously with waterOxidize readily in airGood conductors of electricity
20 Alkali Metals Have one valence electron Will lose this electron very easilywhen electron is lost the metal gains a stable non-reactive noble gas configurationComparison of the Reactivity of the Alkali Metals
21 Group 2: Alkaline Earth Metals Harder, denser, stronger, and have higher melting points than alkali metalsAll are reactive not as reactive as group 1
22 Alkaline Earth MetalsMust lose two electrons to gain a stable configuration
23 Groups 3-12:Transition Metals Not as reactive as Groups 1 and 2Huge variety but all shinyMulti valent…form multiple ionsd-block elementsAlso include: Inner Transition Elements (Rare Earth Elements)Elements LanthanidesElements Actinides
24 Group 17: Halogens Most reactive non-metals Combine easily with metals; especially the alkali metals
25 Halogens 7 valence electrons, one short of a stable octet. Will gain one electron to become stable-1 ionsReaction of chlorine (a halogen) with sodium (an alkali metal)https://www.youtube.com/watch?v=1xT4OFS03jE
27 Hydrogen Most common element in the universe Chemical family by itself because it behaves so differentlyReacts with most other elementsRarely found in a free state in nature1 valence electron
28 The Hindenberg Filled with H Very reactive with oxygen gas He used in blimps todaymuch less reactive than H
29 Group 18: Noble Gases Very low reactivity Filled valence shells: s and p levels in the highest principal energy levels are fullVery stable electron configurationMany uses: signs, weather balloons and the airships (Blimps)
30 The Octet RuleAtoms tend to gain, lose or share electrons in order to acquire a full set of eight valence electrons.Elements on the left (metals) tend to lose valence electrons and form positive ionsElements on the right (nonmetals) tend to gain electrons to become negative ions
32 Periodic TrendsProperties of Elements tend to occur in a predictable wayKnown as a trend, as you move across a period or down a groupKnowing element trends allows us to make predictions about an element’s behavior
33 Periodic Properties Questions we will answer: Properties Atomic Radius DefinitionHow does the property vary across the table?Why?How does it vary down a group?PropertiesAtomic RadiusIonic RadiusElectronegativityIonization Energy
34 Atomic RadiusFor elements that occur as molecules, the atomic radius is half the distance between nuclei of identical atoms.
35 Atomic Radius The atomic radius is a measure of the size of an atom. The larger the radius, the larger is the atom.
36 Trends in Atomic Radius There is a general decrease in atomic radius from left to right, caused by increasing positive charge in the nucleus.Valence electrons are not shielded from the increasing nuclear charge because no additional electrons come between the nucleus and the valence electrons.
37 Trends in Atomic Radius The atomic radius decreases as you move across a periodWhy?Increased nuclear charge pulls the electrons in tighterAdded electrons are in the same principal energy levels
38 Group Trends in Atomic Radius Atomic Radius increases as you move down a groupWhy?The increasing number of electrons are in higher energy levels and instead of pulling the electrons closer to the nucleus we see the …
40 Atomic radius generally increases as you move down a group. The outermost orbital size increases down a group, making the atom larger.
41 Shielding EffectMore inner electrons shield the outer electron from the nucleus and reduce their attraction to the nucleus therefore the overall atomic radius is larger
42 Ionic Radius Atoms can gain or lose electrons to form ions Ion: an atom with a chargeRecall that atoms are neutral in charge,If an electron is lost, then the overall charge is positiveIf an electron is gained the atom becomes negative
43 Positive Ion (Cation) Formation When atoms lose electronsRadius always becomes smaller Because…The loss of a valence electron can leave an empty outer orbital resulting in a small radius.Electrostatic repulsion decreases allowing the electrons to be pulled closer to the radius.
44 Negative Ion (Anion) Formation When atoms gain electronsRadius always increases Why?More electrons mean more electrostatic repulsion resulting in increased diameter.
45 Period Trend for Ionic Radius As you move left to right across a periodthe ionic radius gets smaller for the positive ionsThe ionic radius for the negative ions also decreases
46 Group Trend for Ionic Radius Both positive and negative ions increase in size moving down a group.
48 Ionization Energythe amount of energy need to remove an electron from a specific atom or ion in its ground state in the gas phaseHigh Ionization Energy: atom is holding onto electrons very stronglyLow Ionization Energy: atom is holding electrons less tightly
49 For any element (A) the process of removing an electron can be represented as follows: A + energy -----> A+ + e-What is the periodic trend in ionization energy? Why?
50 Trends for Ionization Energy Generally increases as you move across a periodbecause increased nuclear charge causes an increased hold on the electronsIonization Energy decreases as you move down a groupdue to increasing atomic size
52 Successive Ionization Energies There is an ionization energy for each electron that is removed from an atomAfter the valence electrons are removed Ionization Energies Jump DramaticallyFirst Ionization Energy: removes 1 electronSecond Ionization Energy: removes a second electronThird Ionization Energy: removes a third electron
55 ElectronegativityThe ability of an an atom to attract electrons to itself when it is combined with another atomExpressed in terms of a relative scale: fluorine is assigned a value of 4 and all other elements are calculated relative to this.The units of electronegativity are arbitrary units called Paulings.Noble gases have no values because of few chemical compounds
56 Electronegativity Greater the electronegativity the higher an atom’s ability to pull an electron to itself when it is bonded to another atomWhat are the periodic trends in electronegativity?Why?
57 Trends in Electronegativity Electronegativity Increases as you move across a periodElectronegativity decreases you move down a groupWhere are the elements with highest electronegativity?Where are the elements with lowest electronegativity?