Early Atomic Models. Atoms: the smallest particle of an element that retains the properties of that element.

Transcription

1 Chapter 5

2 Early Atomic Models Atoms: the smallest particle of an element that retains the properties of that element. (Greek: atomos = indivisible) Democritus (Greek teacher in the 4 th century BC) First suggested the idea that atoms existed

3 The Atom 1700 s chemists were able to relate changes to individual atoms Average atom size: Mass = 1 x g Diameter = 1 x 10-8 cm How small is that?100,000,000 copper atoms in a row would = 1 cm in length!

4 Law of Conservation of Mass Definition: mass cannot be created or destroyed only transformed

5 Law of Definite Proportions Definition: a chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source Example: Sodium chloride: NaCl always consists of exactly 39.34% sodium & 60.66%chlorine by mass Water: H 2 O always consists of exactly 11.18% hydrogen & 88.82% oxygen by mass

6 Law of Multiple Proportions Definition: if two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers Examples: CO & CO 2 : 1:1 ratio & a 1:2 ratio H 2 O & H 2 O 2 : 2:1 ratio & a 2:2 ratio

7 John Dalton English school teacher Proposed an explanation for the 3 laws Established in 1808

8 Dalton s Atomic Theory 1. All elements are composed of tiny indivisible particles called atoms. 2. Atoms of the same element are identical. The atoms of any one element are different from those of any other element.

9 3. Atoms of different elements can combine with one another in simple whole number ratios to form compounds. H 2 O C 12 H 22 O 11 NOT H 2.5 O ¾ 4. Chemical reactions occur when atoms are separated, joined or rearranged. Atoms of one element are not changed into atoms of another by chemical means!

10 5. Atoms can not be subdivided.

11 The Structure of the Atom Most of Dalton s Atomic Theory is accepted One major revision includes that idea that atoms are indivisible. There are 3 parts to an atom. 1. electrons 2. protons 3. neutrons

12 Discovery of the Electron Negatively charged subatomic particles J.J. Thomson discovered in 1897 Passed a electric current through gases at low pressures called a Cathode Ray Tube Noticed the surface of the tube directly opposite the cathode glowed. Why? Opposites attract and the electrons were attracted to the positive ends and lights up! Cathode Ray Tube

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14 Cathode Ray Tube

15 Cathode rays are identical regardless of the element Therefore all elements must have electrons! Other important findings: Atoms are electrically neutral, so they must contain a positive charge to cancel it out Since electrons are so small, atoms must contain other particles that account for their mass

16 Robert Millikan ( ) Found quantity of charge in 1 electron (e - ) Also determined the ratio of the charge to the mass of 1 e - Calculated the mass of 1 e - Electrons weigh x10-31 kg

17 J.J. Thomson plum-pudding model e - are spread evenly though out the positive charge of the rest of the atom Ms. Agostine s mint chocolate chip ice cream model

18 Ernest Rutherford (1911) nucleus of the atom is positively charged Gold Foil Experiment

19 Most particles go straight through Positively charged particles deflect off of the positively charged nucleus(~1/8,000) Gold Foil Experiment

20 it was as if you fired a 15-inch [artillery] shell at a piece of tissue paper and it came back and hit you. Nucleus was very small If a nucleus were a marble the atom would be a football field

21 Protons (p + ) Positively charged particles Mass = x kg 1,836 times heavier then an electron

22 Neutrons (n o ) Subatomic particles with no charge Discovered by Sir James Chadwick Mass is nearly the same as a proton Mass = 1.675x10-27 kg

23 Particle Symbol Relative Charge Relative Mass (amu) Actual Mass (kg) Electron e - 1-1/ x10-31 Proton p x10-27 neutron n o x10-27

24 Atomic Number : the number of protons in the nucleus of an atom of an element Atoms are electrically neutral Tells how many electrons there are also! Periodic Table #1 Hydrogen: has 1 p + and 1 e - #6 Carbon: has 6 p + and 6 e -

25 Mass Number total number of protons and neutrons in a nucleus # of neutrons = mass # - atomic # = (# p + + # n o ) - (# p + ) Ex) Beryllium 9 Hyphen notation: The number 9 is the mass number # of p +? # of n o? # of e -?

26 Definition atoms that have the same number of protons but different numbers of neutrons Different types of the same element Ex) Carbon has 3 isotopes Carbon 12 Carbon 13 Carbon 14 Differ by # of n o

27 All have the same # of p + If not, it would be a different element All have 6 protons Carbon 12 Has 6 neutrons Carbon 13 Has 7 neutrons Carbon 14 Has 8 neutrons

28 Hydrogen-1: 1 p + and 0 n o Relative abundance = % Commonly called normal hydrogen Hydrogen-2: 1 p + and 1 n o Relative abundance = 0.015% Commonly called heavy hydrogen or deuterium Hydrogen-3: 1 p + and 2 n o Relative abundance = ~0.00% Commonly called tritium

29 Definition weighted average mass of the atoms in a naturally occurring sample of the element Carbon-12 = % abundant Carbon-13 = 1.11% abundant Carbon-14 = ~ % abundant

30 Formula: Atomic = relative mass # + relative mass # + mass abund. # abund. Repeats for as many isotopes as exist for that element. Units: atomic mass unit (amu): defined as exactly 1/12 the mass of a carbon-12 atom 1 amu = approximately the mass of 1 proton amu s are used so you don t have to use scientific notation when talking about such small masses

31 Sample Problem: Chlorine has 2 isotopes: chlorine-35 which is 75.77% abundant and chlorine-37 which is 24.33% abundant. What is the atomic mass of chlorine?

32 35 Cl = 75.77% abundant = rel. abund. 37 Cl = 24.33% abundant = rel. abund. Atomic mass = = (35 amu x ) + (37 amu x ) = ( amu) + ( amu) = 35.5 amu Compare to value on Periodic Table = amu which rounds to 35.5 amu