Dmitri Mendeleev

• order elements by atomic mass• saw a repeating pattern of properties • Periodic Law – When the elements are

arranged in order of increasing relative mass, certain sets of properties recur periodically

• used pattern to predict properties of undiscovered elements

• where atomic mass order did not fit other properties, he re-ordered by other properties

1834 - 1907

Periodic Table Pattern

H

nm H2O a/b

1 H2

Lim Li2O

b

7 LiH

Bem/nm BeO

a/b

9 BeH2

nm B2O3

a

11 ( BH3)n

Bnm CO2

a

12 CH4

Cnm N2O5

a

14 NH3

Nnm O2

16 H2O

Onm OF2

19 HF

F

left right across a row (period) increasing atomic mass (not relative mass);

organized based on their activity and properties

H

nm H2O a/b

1 H2

Lim Li2O

b

7 LiH

Nam Na2O

b

23 NaH

Bem/nm BeO

a/b

9 BeH2

m MgO

b

24 MgH2

Mg

nm B2O3

a

11 ( BH3)n

B

m Al2O3

a/b

27 (AlH3)

Al

nm CO2

a

12 CH4

C

nm/m SiO2

a

28 SiH4

Si

nm N2O5

a

14 NH3

N

nm P4O10

a

31 PH3

P

nm O2

16 H2O

O

nm SO3

a

32 H2S

Snm Cl2O7

a

35.5 HCl

Cl

nm OF2

19 HF

F

Periodic Table Pattern

up down column; similar properties; how the elements react with H and O

• some gaps were left in table for undiscovered elements; properties of these elements were predicted.

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Mendeleev's Predictions for Ekasilicon (Germanium)

Property Silicon’s Props

Tin’s Props

Predicted Value

Measured Value

Atomic Mass

28 118 72 72.6

Color Grey White metal

Grey Grey- White

Density 2.32 7.28 5.5 5.4

Reaction w/ Acid &

Base

Resists Acid, Reacts Base

Reacts Acid,

Resists Base

Resists Both

Resists Both

Oxide SiO2 SnO2 Eks1O2 GeO2

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72.61

Ge

14

28.09

Si

Sn 50

118.71

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PERIODIC TABLE

Metals Non-metals

Metalloids

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PROPERTIES OF METALS & NON-METALS

Metals Non-metals

• Mostly solid • Can be solid, liquid or gas

• Have shiny appearance • Have dull appearance

• Good conductors of heat & electricity

• Poor conductors of heat & electricity

• Malleable & ductile • Brittle (if solid)

• Lose electrons • Gain or share electrons

Metals Non-metals

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METALLOIDS

Metalloids are elements that possess some properties of metals and some of non-metals.

The most important metalloids are silicon (Si) and germanium (Ge) which are used extensively in computer chips.

Metalloids

Properties of Siliconshiny

conducts electricitybrittle

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PERIODIC TABLE

Metallic character increases going down a groupMetallic character decreases going across a period.

Cs

Fr

Most metallic elements F

Least metallic element

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PERIODIC TABLE

Seven elements exist as diatomic molecules. All others exist as monatomic (single atom).

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PERIODS & GROUPS

The periodic table is composed of periods (rows) and groups or families (columns).

Elements in the same family have similar properties, and are commonly referred to by their traditional names.

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PERIODS & GROUPS

Elements in groups 1-2 and 13-18 are referred to as main-group or representative groups.

Alkali metals are soft metals that are very reactive. They often react explosively with other elements.

Noble gases are un-reactive gases that are commonly used in light bulbs.

Halogens are the most reactive nonmetals, and occur in nature only as compounds.

Group 2 elements are called alkaline-earth metals. These metals are less reactive than alkali metals.

The group of metals in between the main group elements are called transition metals.

Important Groups - Hydrogen• nonmetal• colorless, diatomic gas

– very low melting point & density

• reacts with nonmetals to form molecular compounds– HCl is acidic gas and H2O is a liquid

• reacts with metals to form hydrides (negative ion of hydrogen = H-)– metal hydrides react with water to form H2

(Metal-H + H2O H2 + metal-OH)

(NaH + HOH H2 + NaOH)

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Important Groups – IA, Alkali Metals

• hydrogen usually placed here, though it doesn’t belong

• soft, low melting points,low density

• flame tests Li = red, Na = yellow, K = violet

• very reactive, never find uncombined in nature

• tend to form water soluble compounds

lithium

sodium

potassium

rubidium

cesium

Important Groups – IIA, Alkali Earth Metals

• harder, higher melting, and denser than alkali metals

• flame tests Ca = red, Sr = red, Ba = yellow-green

• reactive, but less than corresponding alkali metal

• form stable, insoluble oxides from which they are normally extracted

• oxides are basic = alkaline earth

• reactivity with water to form H2,

magnesium

calcium

beryllium

strontium

barium

Important Groups – VIIA, Halogens

• nonmetals

• F2 & Cl2 gases; Br2 liquid; I2 solid

• all diatomic• very reactive• react with metals to form ionic

compounds• HX (X = halogen) are all acids

– HF weak < HCl < HBr < HI

bromine

iodine

chlorine

fluorine

Important Groups – VIIIA, Noble Gases

• all exist as gases at room temperature, – very low melting and

boiling points

• very unreactive, practically inert

• very hard to remove electron or add an electron

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ATOMICSTRUCTURE

The general designation for an atom is shown below:

Atomic number (Z) = # of protons

Mass number (A) = # of p+ + # of n0

# of n0 = A - Z

Charged Atoms• The number of protons determines the

element!– all sodium atoms have 11 protons in the nucleus

• In a chemical change, the number of protons in the nucleus of the atom doesn’t change!– no transmutation during a chemical change!!– during radioactive and nuclear changes, atoms do

transmute

• Atoms in a compound are often electrically charged, these are called ions

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5.1

ELEMENTS& IONS

An ion (charged particle) can be produced when an atom gains or loses one or more electrons.

A cation (+ ion) is formed when a neutral atom loses an electron

Metals form cations

Cations are named the same as the parent atom

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5.1

ELEMENTS& IONS

An anion (- ion) is formed when a neutral atom gains an electron

Non-metals form anions

Anions are named by using the root of the parent atom’s name and changing the ending to –ide.

Chlorine changes to chloride

Ion Charge & the Periodic Table

• Every element wants to be like the noble gases• wants 0 or 8 valence electrons• valence e- (outer electron shell used for bonding)• He has zero valence e-

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elements of Group IIA have a

+2 charge

elements of Group IA have a +1

charge

elements of Group VA have a

-3 charge

elements of Group VIA have a

-2 charge

elements of Group VIIA have a

-1 charge

ELEMENTS& IONS

For main-group elements, the charge of ions are very characteristic of the group numbers.

Structure of the Nucleus

• Frederick Soddy discovered that the same element could have atoms with different masses, which he called isotopes– there are 2 isotopes of chlorine found in nature, one

that has a mass of about 35 amu and another that weighs about 37 amu

• The observed mass is a weighted average of the weights of all the naturally occurring atoms– the atomic mass of chlorine is 35.45 amu

(1877-1956)

Isotopes

• Atomic mass - The average mass for a given element• all isotopes of an element have the same number of

protons• isotopes of an element have different masses• isotopes of an element have different numbers of

neutrons• isotopes are identified by their mass numbers

– protons + neutrons

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ISOTOPES

Atoms of the same element that possess a different number of neutrons are called isotopes.

Isotopes of an element have the same atomic number (Z), but a different mass number (A).

The 3 isotopes of Hydrogen

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ISOTOPES &ATOMIC MASS

The mass of an atom is measured relative to the mass of a chosen standard (carbon-12 atom), and is expressed in atomic mass units (amu).

The average atomic mass of an element is the mass of that element’s natural occurring isotopes weighted according to their abundance.

Therefore the atomic mass of an element is closest to the mass of its most abundant isotope.

4.8 Isotopes: Natural AbundanceIsotopes of neon  Naturally occurring neon contains three different isotopes: Ne-20 (with 10 protons and 10 neutrons), Ne-21 (with 10 protons and 11 neutrons), and Ne-22 (with 10 protons and 12 neutrons).

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Example 1:

Determine the number of protons, electrons and neutrons in a chlorine atom .

3517Cl

A = 35

Z = 17

# of protons = 17 (Z)

# of electrons = 17 (= p+)

# of neutrons = 18 (35 - 17)

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Example 2:

Which two of the following are isotopes of each other?

410 410 412 412186 185 183 185X Y Z R

Isotopes of an element have the same atomic number, but a different mass number

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Example 3:

Based on the information below, which is the most abundant isotope of boron (atomic mass = 10.8 amu) ?

Atomic mass of an element is closer to the mass of the more abundant isotope

Isotope 10B 11B

Mass (amu) 10.0 11.0

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(106.91) (0.5184) = 55.42 amu

(108.90) (0.4816) = 52.45 amu

107.87 amu

IsotopeMass

(amu)Abundance

(%)

107Ag 106.91 51.84

109Ag 108.90 48.16

CALCULATING MASSFROM ISOTOPIC DATA

Atomic mass Abundance Mass of Abundance Mass ofx + x

of an element of isotope 1 isotope 1 of isotope 2 isotope 2

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Mass Number is Not the Sameas Atomic Mass

• the atomic mass is an experimental number determined from all naturally occurring isotopes

• the mass number refers to the number of protons + neutrons in one isotope– natural or man-made

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THE END