Chapter 2 Atoms and Elements

atom- smallest identifiable unit of an element

element- a substance that cannot be broken down into simpler substances

-there are about 91 naturally occurring elements

-scientists have made over 20 elements

Modern Atomic Theory and Laws That Led to It

1) Law of conservation of mass- matter is neither created nor destroyed

-when a reaction is carried out, the mass of the reactants must equal the mass of the products

ex- 2Na + Cℓ2 2NaCℓ

7.7g 11.9g 19.6g

19.6g

2) Law of Definite Proportions- all samples of a given compound have the same proportions of their constituent elements

ex- 18.0g H20 results in 16.0g O2 and 2.0g H2

mass ratio = 16.0g O2/2.0g H2 = 8 O2:1 H2

-this is true for all samples of water

Page 49 Ex 2.1 and For practice 2.1

3) Law of Multiple Proportions- when two elements (A and B) form two different compounds, the mass of element B that combines with 1g of element A can be expressed as a ratio of small whole numbers

Page 50

Ex- 2.2

For practice 2.2

John Dalton’s Atomic Theory

1. All elements are composed of atoms

2. All atoms of the same element have the same properties and are identical

3. Atoms combine in simple whole number ratios with other atoms to form compounds

4. Atoms of one element cannot be changed into atoms of another element

-Scientists began to think that atoms were composed of smaller particles

Discovery of Electron

-J.J. Thomson (1856-1940)

-used a cathode ray tube

-believed that the cathode ray was composed of tiny particles with an electrical charge

-these are electrons

-rays travel from – charged electrode (cathode) to + electrode (anode)

Discovery of Charge of Electron

-Robert Millikan (1868-1953)

-said electrons have a negative charge

Plum Pudding Model

-it was known that – charged particles attract + charged particles

-it was also known that atoms were neutral

-so there must be a + charged particle in the atom

-Thomson proposed that – charged electrons were small particles held within a + charged sphere

-protons discovered by Eugen Goldstein (1850 – 1930)

Ernest Rutherford (1871-1937)

-performed experiments that allowed him to conclude that the atom must have a positive mass in a much smaller space than proposed

Nuclear Theory

1) most of atom’s mass and all of its + charge are in a small core called the nucleus

2) most of volume is empty space through which negative particles move

3) there are as many negatively charged particles as there are positive (protons) so the atom is neutral

-There were still some parts of atom missing

James Chadwick (1891-1974)

-discovered that missing mass was neutrons

Subatomic Particles and the Atom

protons (p+) – positively charged, found in nucleus, mass = 1.67 x 10-24g

neutrons (n0) – neutrally charged, no charge, found in the nucleus, mass = 1.67 x 10-24g

electrons (e-) – negatively charged, found surrounding nucleus in clouds or energy levels,

mass = 9.11 x 10-28g

**most of mass in the nucleus of an atom

atomic number- # of protons in an element, defines the element, the smaller # on the periodic table

-since atoms are neutral:

# of p+ = # of e-

-so the atomic # also tells you the # of e-

How many protons in the following?

argon uranium iron lithium

18 92 26 3

*each element also has that # of e-

mass number- sum of the # of p+ and n0, larger # on the periodic table rounded to a whole #

How do you find the # of n0?

# n0 = mass # - atomic #

How many neutrons in the following?

argon uranium iron lithium

22 146 30 4

Can be written this way:

Ar-40 U-238 Fe-56 Li-7

atomic mass- larger # on the table not rounded, average mass of all the isotopes of that element

isotopes- atoms with the same # of p+ and e- but different # of n0

ex-

Ne-20 Ne-21 Ne-22

p+ = 10 p+ = 10 p+ = 10

e- = 10 e- = 10 e- = 10

n0 = 10 n0 = 11 n0 = 12

Page 59 Ex 2.3 and For Practice 2.3

natural abundance- the relative % of an isotope with respect to other isotopes of the same element

-back to Ne isotope

Which isotope of Ne would have the highest relative abundance?

Ne-20, because it is closest to the atomic mass

-Relative Abundance problems

Modern Periodic Table

-grew from the work of Dmitri Mendeleev

-arranged elements according to increasing mass

-called periodic law- when elements are arranged in order of increasing mass/atomic #, similar properties will recur

periods- horizontal rows

groups/families- vertical columns, indentified by a # and a letter

Groups of the Periodic Table

Alkali Metals- Group 1A elements

Alkaline Earth Metals- Group 2A elements

Transition Metals- Group B elements, bridge

Inner Transition Metals- two rows at the bottom

non-metals- upper and lower right side

Metalloids- staircase- starting at boron

Halogens- Group 7A

Noble Gases- Group 8A or Group 0

-also called inert gases- they do not react to form compounds

Main group or Representative elements

-Groups 1A to 8A minus the Transition metals

Properties of Metals

1) conduct electricity

2) ductile (can be drawn into wires)

3) shiny/lustrous

4) malleable (can be pounded into sheets)

-Most elements are solids

Gases = H, O, N, F, Cℓ, He Ne, Ar, Kr, Xe, Rn

Liquids = Hg, Br (Ga, Fr, Cs)