The Periodic Table of Elements

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Periodic Table• Something “periodic” occurs at regular

or generally predictable intervals • Periodic law - physical and chemical

properties of the elements are periodic functions of their atomic numbers

• Periodic Table of Elements – a table of the elements, arranged by atomic number, that shows the patterns in their properties; based on the periodic law Can you think of anything that is periodic? 2

Element• A pure substance made up of one kind

of atom that cannot be broken down into simpler substances by physical or chemical means

• 90 occur naturally on earth• 25 were synthesized (made) by

scientists• The Element Song

http://www.privatehand.com/flash/elements.html

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Dmitri Mendeleev• In the 1860’s he

devised a periodic table where the elements were ordered by their atomic masses

• He did this by grouping elements together according to their similarities

• Draft of Mendeleev's Periodic Table

Image taken from: http://jscms.jrn.columbia.edu/cns/2006-04-18/fido-luxuriantflowinghair/mendeleev/ 4

Why do you think there are question marks here?Image taken from: http://www.chemsoc.org/networks/learnnet/periodictable/post16/develop/mendeleev.htm

Mendeleev’s Published Periodic Table of Elements

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Mendeleev’s Predictions• Although Mendeleev’s Periodic Table of

Elements had missing elements or “gaps,” he was able to predict the characteristics of these missing elements because of Periodic Law.Date Predicted

1871 Date Discovered

1886

Atomic Mass 72 Atomic Mass 72.6

Density 5.5 g/cm3 Density 5.47 g/cm3

Bonding Power

4 Bonding Power

4

Color Dark Gray Color Grayish White

“Ekasilicon” GermaniumNotice how

Mendeleev’s predictions

(orange column) were very

accurate when compared to Germanium’s

actual characteristics (green column)

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Henry Moseley• In 1914, his work led

to a revision of the periodic table by rearranging the elements by their atomic numbers

• He concluded that the number of protons in an atom is its atomic number

Image taken from: http://dewey.library.upenn.edu/sceti/smith/ 7

3 Classes of ElementsUsing this as a guide, color code your periodic table to

show the classes.

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MetalsLocation • Found on the left of the

zigzag line/staircase on the periodic table (exception Hydrogen)

Chemical Properties• Have few electrons in

their outer energy level, thus lose electrons easily

Physical Properties• Ductile, good conductors,

malleable, shiny, most are solid @ room temperature

11Na

22.990

79Au

196.967

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Reference-Metals’ Chemical Properties

Notice: only 1 electron in outer

level1s22s22p63s1

Notice: only 2

electrons in outer level

1s22s2

++

+

+

+++

+++

-

-

-

-

--

-

-+

-

-

-+

+++ --

-

-

11Na

22.990

4Be

9.012

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Metals’ Physical Properties

• Good conductor- electrons (electricity) flow easily through the substance

• Malleable- able to be hammered or pressed out of shape without breaking

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Non-MetalsLocation • Found to right of the

zigzag line/staircase Chemical Properties • Most have almost full outer

energy levels, so tend to gain electrons; some completely full

Physical Properties • Not ductile or malleable, not

shiny, poor conductors, most are solid, but some are gas at room temperature

16S

32.066

17Cl

35.453

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MetalloidsLocation

• Border the zigzag line/staircase on the periodic tableChemical Properties

• Most atoms have ½ (≈) complete set of electrons in outer levelPhysical Properties

• have properties of both metals and non-metals

• B, Si, Ge, As, Sb, Te, At

5B

10.811

14Si

28.086

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Using the Periodic Table• The boxes that make up the periodic

table contain a significant amount of information Atomic Number

(number of protons)Element Symbol(capital letter or a capital and lower case)

Atomic Mass (weighted averages)

8O

Oxygen15.999 Element Name

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• To understand this information, it is necessary to refer to the periodic table’s key(s)

Important Features of the Periodic Table

• Period- each horizontal row of elements on the periodic table

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Period (Series) Properties• Every element belongs to a period• Seven periods on a periodic table

(numbered from the top down) • Period number = quantum number “n”

= highest energy level of atoms where valence electrons are found

– Carbon: 1s22s22p2 2nd period 131– 40Zr [Kr] 5s24d2 5th period

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Period (Series) Properties• Atomic numbers and atomic masses

increase as you move from the left to the right in a period

• Each element in a specific period has that respective number of levels– Ex

• Period 1 elements = level 1 valence electrons

• Period 2 = level 2 • Period 3 = level 3

• Etc…

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• Group- column of elements on the periodic table

Important Features of the Periodic Table

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Group (Family) Properties• Eight main groups (numbered from left

to right)• Not all elements belong to a group or

family• Atomic numbers and atomic masses

increase as you move from the top down in a group (family)

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Group (Family) Properties• Elements in a group have similar

properties and electron configurations• In an electron configuration, the

NUMBER of valence electrons in the highest level is that element’s group, or family

108• Ex 33As [Ar] 4s23d104p3

Group V- Nitrogen Family 20

GROUP FAMILY NAME EXAMPLEI Alkali Metals NaII Alkaline Earth Metals CaIII Aluminum Family AlIV Carbon Family CV Nitrogen Family NVI Chalcogen Family OVII Halogen Family FVIII Noble or Inert Gas Family Ne

Group #’s and Family Names-

Reference

Group (Family) NamesLabel your periodic tableAlkali

Metals Alkaline Earth MetalsTransition

Metalsnot a group

Boron GroupCarbon

GroupNitrogen

Group HalogensNobl

e Gase

s

22

Chalcogen Group

Let’s Do It!!!!

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• Identify the element:• Group 2 Period 2 = • Group 7 Period 6 =• Group 6 Period 4 =• Group 1 Period 4 =• Group 4 Period 1 =

Let’s Do It!!!!

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• Identify the element:• Group 2 Period 2 = Be (Beryllium)• Group 7 Period 6 = At (Astatine)• Group 6 Period 4 = Se (Selenium)• Group 1 Period 4 = K• Group 4 Period 1 = Nada, nothing,

zero, emptiness

3 Classes of Elements

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Transition Metals• Do not belong to a group and all are

metals• Element where electrons are being

placed into a “d” sublevel and “d” is last in electron configuration

• Ex: 23V [Ar] 4s23d3 • Similar to Group II elements but some

different properties because of the “d” electrons

• They contain two dots in their dot structure

Rare Earth (Inner Transition) Metals

• Do not belong to a group and all are metals• Element where electrons are being placed into a “f” sublevel and “f” is last in electron configuration

Example: 92U [Rn] 7s25f4

• Similar to Group II elements but also some different properties because of the “f” electrons•All contain two dots in their dot structure

Summary• An element…

– can be in a group OR

– can be a transition element OR

– can be a rare earth element

• It may NOT belong to more than 1 of these 3 categories

ReviewPeriods and Groups video

10min• Number your paper from 1-5 and answer the following questions. Two will be cumulative review!

• 1. Which of the following explains the difference between an atom and an ion?A. Ions have different numbers of

protons and neutronsB. Ions have different numbers of

protons and electronsC. Ions have different numbers of

neutrons and electrons

– b. – c. – d.

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Review• B• 2. What is the electron configuration of

Calcium?– a. [Ar]4s2 – b. [He]2s1

– c. [Ne]3s2

– d. [K]4s2

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Review• A • 3. What is true of elements in the same

group?A. They have similar electron

configurationsB. They have similar propertiesC. They are members of the same

chemical familyD. All of the above

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Review• D• 4. Which of the following describes the

element chromiumA. In group IB. In group IIC. A transition elementD. In group IV

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Review• C • 5. How many valence electrons do

transition metals have?– A. 1– B. 2– C. 3– D. 4

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Review• B….two dots in their Lewis dot

structure

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Periodic Trend #1: Metals vs. Non-metals

• Metallic- able to lose electrons• DOWN GROUPS, elements become MORE

metallic– Electrons are further from pull of positive

protons so lost easier = conduct electricity

• ACROSS PERIODS, elements become LESS metallic– More protons so more attraction– Gain more electrons and becoming more

stable and less likely to lose electrons

Reference-Periodic Trend #1: Metallic Properties

Video

Periodic Trend #2: Atom Size• 2 things affect the size of atoms:

– Increasing the number of electrons increases the size (the more electrons, the bigger the atom)

– Increasing the number of protons increases the attraction between the nucleus and electrons and decreases the size

• These two statements contradict each other so which one is correct?

• ACROSS PERIODS atoms DECREASE in size as electrons are being added to the same energy level (reason #2 is dominant over #1)

ATOM Na Mg Al S P S Cl Ar

SIZE (radius in Å) 1.54 1.36 1.18 1.11 1.06 1.02 .99 .98

Periodic Trend #2: Atom Size

• DOWN GROUPS atoms INCREASE in size as the electrons go into a higher energy level (get farther away from the nucleus) (Reason #1 above is dominant over #2)

ATOM SIZE (radius in Å) H .32 Li 1.23

Na 1.54 K 2.03 Rb 2.16 Cs 2.35

Reference- Periodic Trend #2: Atom Size

Reference- Periodic Trend #2: Atom Size

Video

Ions

• Ion- an atom with a charge due to different numbers of protons and electrons

• Ex.) – Na atom 11 protons, 11 electrons =

neutral– Na ion 11 protons, 10 electrons = +1

charge• With new electron structures, ions have

different properties than the atoms from which they were formed

• For an atom to “lose” an electron, a certain amount of energy needs to be absorbed

Periodic Trend #3Ionization Energy (I.E.)

• Ionization energy- energy needed to remove an electron from an atom and form a positive ion

• Generic formula X(atom) + ionization energyX+1(ion) +

electron• Specific example Na + 119 Kcal/mole Na+1 +

electron • The ion is positive because it has 1 less

electron than protons– Positive ions = “cations” – Smaller radius than original atoms

• ACROSS PERIODS ionization energy generally INCREASES– Increased attraction between the +

nucleus and the electrons, hence a higher ionization energy

ATOM: Na Mg Al Si P S Cl Ar

I.E. (kcal/mole): 119 176 138 188 242 239 299 363

Reference- Periodic Trend #3Ionization Energy (I.E.)

• DOWN GROUPS ionization energy generally DECREASES

• Electrons in higher energy levels and farther away from the positive nucleus so less energy is required to remove them

ATOM I.E. (kcal/mole) H 314 Li 124

Na 119 K 100 Rb 96 Cs 90

Cesium’s outer electron is much further away from the nucleus, therefore it is easier to remove

Reference-Periodic Trend #3Ionization Energy (I.E.)

Reference- Periodic Trend #3Ionization Energy (I.E.)

Video

• The lower the ionization energy, the more chemically active the element will be – It will lose electrons and react or

“bond” with other elements more readily

– Metals have low ionization energies • Few electrons in their highest energy

level• Electrons are weakly held and easily

removed

Periodic Trend #3Ionization Energy (I.E.)

• A metal will lose enough electrons until a stable electron structure, (usually a noble gas structure) is obtained:– Metals in Group I with s1 configurations

will lose 1 electron and form +1 ions– Metals in Group II with s2

configurations will lose 2 electrons and form +2 ions

– Metals in Group III with s2p1 configurations will lose 3 electrons and form +3 ions

Periodic Trend #3Ionization Energy (I.E.)

Reference- Electron Configuration Comparison

Atom config. of atom Ion config of ionNa 1s22s22p63s1 Na+1

1s22s22p6

Mg 1s22s22p63s2 Mg+2 1s22s22p6

Al 1s22s22p63s23p1 Al+3 1s22s22p6

not stable stable

I.E. (cont.)• Most transition and rare earth elements

vary in the number of electrons they lose so their charge will vary (can lose electrons from several sublevels)

• Some transition elements, like silver, always lose the same number of electrons and always have the same charge

• Non-metallic elements do NOT form positive ions as they would have to lose too many electrons to form a positive ion with an electron structure like a noble gas

Periodic Trend #4: Electron Affinity (E.A)

• Electron affinity- energy released when an electron is gained by an atom

General Formula X (atom) + electron ---> X-1 (ion) +

energy Specific Formula Cl + electron ---> Cl-1 + 83.4

kcal/mole • This ion is negative because there is 1 more

electron than proton• A negative ion = anion• Larger than the original atoms

• ACROSS PERIODS electron affinity INCREASES to a maximum in Group VII then is a minimum in Group VIII

ATOM: Na -> Cl Ar E.A.: 12.6 -> 83.4 -8.5

(kcal/mole)• DOWN GROUPS electron affinity usually

DECREASES due to the positive nucleus being less attracted to the electrons further away

Reference- Periodic Trend #4:

Electron Affinity (E.A)

• Core elements in Group VIII (s2p6): already stable electron structures and don’t want any more

• Group VII (s2p5): benefit greatly by getting 1 more (s2p6) so the ion stable & form -1 ions

• Group VI: (s2p4) gain 2 e- (s2p6) & form -2 ions• Group V: (s2p3) gain 3 e- (s2p6) & form -3 ions • Metal elements do NOT form negative ions as

they would have to gain too many electrons to have an electron structure like an inert gas (core atom)

Periodic Trend #4: Electron Affinity (E.A)

Atom e configuration Ion e configuration

Cl 1s22s22p63s23p5 Cl-1 1s22s22p63s23p6

S 1s22s22p63s23p4 S-2 1s22s22p63s23p6

P 1s22s22p63s23p3 P-3 1s22s22p63s23p6

not stable stable

Reference- Periodic Trend #4:

Electron Affinity (E.A)

Reference-Periodic Trend #4: Electron Affinity (E.A)

Video

Ion Summary-ReferenceGROUP I II III IV V VI VII VIII

Electron Configuration Ends in:

s1 s2 s2p1 s2p2 s2p3 s2p4 s2p5 s2p6

Electrons Lost  1 2  3  - - - - -

Electrons Gained - - - - 3  2  1  -

Ion Formed

+1  +2 +3 - -3 -2 -1 -

Essentially the Group Number tells you how many valence electrons there are!!!*Note – there are some exceptions to this chart!

Reference

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Exceptions

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Exceptionalelements

• The elements in this table can form more than one type of ion

• When naming these compounds, the type of ion is expressed in the name with a roman numeral• These are all transition

elements and they can form a variety of ions!

Relating IE, EA and Metallic Properties

*Note – there are some exceptions

• High in metallic properties = low ionization energy = low electron affinity

• Low in metallic properties = high ionization energy = high electron affinity

ReviewIons video 10min

• Number your paper from 1-5 and answer the following questions. Two will be cumulative review!

• 1. Which of these is a metal?– a. S– b. Be– c. F– d. O

59

Review• B• 2. Which of the following describes the

element Tin (Sn)A. In group IB. In group IIC. In group IVD. A transition element

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Review• C• 3. Which is more metallic?

– a. Li– b. Na– c. Cs– d. Fr

61

Review• D• 4. Which of these has the lowest

ionization energy – a. Li– b. Na– c. Cs– d. Fr

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Review• D• 5. Which of these has the lowest

electron affinity– a. Li– b. Na– c. Cs– d. Fr

63

Review• D

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In a chemical reaction…… 1) new substances with new properties

are formed2) no change in mass

Review: “Ions”• The Octet Rule

– Atoms want stable electron configurations so they must achieve a noble gas configuration with eight valence electrons (s2p6)

– Some very small atoms want to be like helium which only has two valence electrons (s2) therefore, they follow the “duet rule”

– In order to follow the octet (or duet) rule, atoms need to lose, gain, or share electrons and in doing so can form ions

Bonding• Recall that a compound is

a combination of two or more elements H2O

• Chemical formula- tells what elements a compound contains and the exact number of the atoms of each element

• Water contains H for the element hydrogen and O for the element oxygen

Formulas• Numbers written before a

symbol tells you how many of those compound units you have 2 H2O = H2O + H2O

• Subscripts written after a symbol tells how many atoms of that element are in a unit of the compound (H2O contains two hydrogen atoms)

• No subscript = only one atom of that element (H2O contains one oxygen atom)

•

Ionic Bonding• When cations and anions bond, a

chemical reaction takes place• Consider the following chemical reaction:

Na + Cl NaCl

#1 Cation StepNa + small energy Na+1 + e-1

• A small amount of energy is absorbed (ionization energy) to release an electron

• The smaller the energy needed, the easier it is for an atom to release an electron

• Metals have low ionization energies

#2 Anion StepCl + e-1 Cl-1 + large energy

• Non-metals receive electrons when available and form negative ions

• A large amount of energy is released (electron affinity)

• The larger the electron affinity, the more likely this reaction takes place

#3 Ionic Compound Formation

Na+1 + Cl-1 Na+1Cl-1 + small energy

•Since opposite charges attract, the ions bond to form a compound•A small amount of energy is released when ions bond and together they are more stable•The resulting compound is called an ionic compound

Reference-Overall Steps to Ionic Bonding

Ex.) Sodium bonding with chlorine1. Na + small energy (I.E) Na+1 + e-1

2. Cl + e-1 Cl-1 + large energy (E.A)3. Na+1 + Cl-1 Na+1Cl-1 + small energyOverall reaction (cancel like terms THAT

ARE ON BOTH SIDES): Na + Cl NaCl + large energy and the large amount of energy given off shows that the elements of K2O are much

more stable together

Reference- Steps to Ionic Bonding

Ex.) potassium bonding with oxygen1. 2K + small energy (I.E) 2K+1 + 2e-1

2. O + 2e-1 O-2 + large energy (E.A)3. 2K+1 + O-2 K2

+1O-2 + small energy

Overall Reaction2K + O K2O + large energy

and the large amount of energy given off

shows that the elements of K2O are much more stable together

Ionic Compounds• Electrons are essentially transferred from

the metal atom to the nonmetal • Properties of the compound are different

from those of the elements which now have different electron structures

• We write ionic compounds as empirical formulas- the smallest whole number ratio between the elements in the compound

• Ionic compounds must be neutral so the overall positive charge must equal the overall negative charge

• A calcium ion has a charge of 2+ and a fluoride ion has a charge of 1 −

• In this case you need to have two fluoride ions for every calcium ion in order for the charges to cancel CaF2

• +2 -1 -1 = 0 neutral

Compounds Are Neutral

• sodium chloride (+1) (-1)

NaCl

• magnesium oxide

(+2) (-2)MgO

• indium selenide (+3) (-2)

In2Se3

Reference-Ionic Compounds

• Ionic compound steps:• 1. Write the symbol of the

ion that has the positive charge (metal)

• 2. Write the symbol of the element with the negative charge (nonmetal)

• 2. CRISS CROSS the numbers of the charges and make them

the subscripts in the empirical formula

• 1. boron iodide

• 2. calcium bromide

• 3. cesium phosphide

Let’s Do It!!!!

C CR R I O N I CS SS S

• 1. boron iodide (+3) (-1) BI3

• 2. calcium bromide (+2) (-1)

CaBr2

• 3. cesium phosphide (+1) (-3)

Cs3P

Let’s Do It!!!!

C CR R I O N I CS SS S

Let’s put on our thinking caps!

• Compound between K and Mg?

• Compound between Cl and O?

• Compound between Cu and S?

Naming Ionic CompoundsLet’s put on our thinking

caps!• Compound between K and Mg?– None, two positives

• Compound between Cl and O? - none, two negatives

• Compound between Cu and S? - There are two!- One with copper (I) and one with copper (II) Cu2S CuS

Writing NamesIonic Compounds

• You can name a binary ionic compound from its formula by using these rules

• Write the name of the positive ion (metal)

• Write the root name of the negative ion second (nonmetal) which is the first part of the element’s name

• Change the last syllable of the nonmetal to -ide

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Writing NamesIonic Compounds

• Subscripts do not become part of the name for ionic compounds

• However, subscripts can be used to help determine the charges of the ions that make it

• Reverse the CRISS CROSS• Cu2S = Cu+1 + S-2

• So the form of the transition metal Copper in this case is Cu(I)

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Ionic Compounds with Complex Ions

• Not all ionic compounds are binary like NaCl and made of only two elements

• Baking soda has the formula NaHCO3

• NaHCO3 is an example of an ionic compound that is not binary because it has more than two elements

84

Polyatomic Ions• These ionic

compounds like baking soda are made of more than two elements

• They have ions made of more than one element

• A polyatomic ion is a positively or negatively charged, covalentlybonded group of two or more elements

85

NH4+1

Polyatomic Ions• To write formulas for

these compounds, follow the rules for binary ionic compounds, with one addition: when more than one polyatomic ion is needed, write parentheses around the polyatomic ion before adding the subscript Mg3(PO4)2

86

NH4+1

ReviewWriting Ionic Compounds

video 10min• Number your paper from 1-5 and

answer the following questions. Two will be cumulative review!

• 1. Which of these is a metal?– a. S– b. Be– c. F– d. O

87

Review• C• 2. What kind of ion would iron (Fe)

make?A. -1B. +2C. +3D. -8

88

Review• B Metals LOSE electrons to become

more stable (s2p6)and Fe is a transition element with two valence electrons

• 3. How many electrons would S have to gain to have a stable noble gas configuration?

A. 2B. 3C. 5D. 8 89

Review• A Nonmetals GAIN electrons to become

more stable (s2p6)• 4. How many electrons does Ca have to

lose to have a stable noble gas configuration?

A. 1B. 2C. 3D. 8

90

Review• B Metals LOSE electrons to become

more stable s2p6 [Ar]• 5. Will an ionic compound form

between K and Ne?– A. Yes– b. No– c. Sometimes

91

Review• B

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