Presentation is loading. Please wait.

Presentation is loading. Please wait.

Fundamental Chemical Laws

Published byJason Poole Modified over 8 years ago

Similar presentations

Presentation on theme: "Fundamental Chemical Laws"— Presentation transcript:

1 Fundamental Chemical Laws

2 Law of Conservation of Mass
Lavoisier “Father of Modern Chemistry” In every chemical operation an equal amount of matter exists before and after the operation. Mass is conserved, the total mass after the chemical operation must be the same as that before.

3 Law of Definite Proportions
Joseph Proust In a given chemical compound, the proportions by mass of the elements that compose it are fixed, regardless of the source of the compound. The ratio of elements in a compound is fixed regardless of the source of the compound. Water is made up of 11.1% by mass of hydrogen and 88.9% oxygen.

4 Avogadro’s Hypothesis
Equal volumes of different gases (at the same temperature and pressure) contain equal numbers of particles 2 volumes of hydrogen + 1 volume of oxygen  2 volumes of water vapor can be expressed as 2H2 + O2  2H2O While at this time there was no direct evidence to show that hydrogen and oxygen gas were H2 and O2, 50 years later this was proven to be the case.

5 Dalton’s Atomic Theory
Elements are made of tiny particles called atoms. All atoms of a given element are identical. The atoms of a given element are different from those of any other element. Atoms of one element can combine with atoms of other elements to form compounds. A given compound always has the same relative number of types of atoms. Atoms cannot be created, nor divided into smaller particles, nor destroyed in the chemical process. A chemical reaction simply changes the way atoms are grouped together.

6 Evidence for sub-atomic particles
1897: J.J. Thomsen: Cathode Ray Tube Evidence for electrons: Bent a stream of rays originating from the negative electrode (cathode). Stream of particles with mass & negative charge. 1909: Ernest Rutherford: Gold Foil Evidence for protons & nucleus: Alpha particles deflected passing through gold foil 1932: James Chadwick: Beryllium Evidence for neutrons: Alpha particles caused beryllium to emit rays that could pass through lead but not be deflected

7 Millikan’s Oil Drop Experiment
Robert Millikan’s oil drop experiment calculated the charge/mass ratio of the electron, and combining Thompson’s results the mass of the electron was calculated to be 9.10 x g. (actual mass of the electron x g) There must be a positive species which counters the electron charge.

8 Radioactivity Henri Becquerel in 1896 discovered high-energy radiation was spontaneously emitted from uranium. Later Marie Curie and her husband Pierre further investigated this spontaneous emission of radiation which was termed radioactivity.

9 The Structure of the Atom
J.J. Thompson, realized that electrons were sub-atomic particles, and presented his theory of the model of the atom. The “PLUM-PUDDING” model

10 Model of the Atom Since the times of Rutherford, many more subatomic particles have been discovered. However, for chemists three sub-atomic particles are all that we need to focus on – ELECTRON, PROTON, NEUTRON. Electrons are –1, protons +1 and neutrons are neutral. Atoms have an equal number of electrons and protons they are electrically neutral.

11 Protons and neutrons make up the heavy, positive core, the NUCLEUS which occupies a small volume of the atom.

12 Isotopes 11B 10B Atoms of the same element but different mass number.
Boron-10 (10B) has 5 p and 5 n Boron-11 (11B) has 5 p and 6 n 10B 11B

13 Two Isotopes of Sodium.

14 Masses of Particles Relative Isotopic Mass
Chemists as early as John Dalton, two centuries ago, used experimental data to determine the weight of different atoms relative to one another. Dalton estimated relative atomic weights based on a value of one unit for the hydrogen atom. In 1961, it was decided that the most common isotope of 12C would be used as the reference standard. On this scale, the 12C isotope is given a relative mass of exactly 12 units.

15 Relative Isotopic Mass cont…
“The relative isotopic mass (Ir) of an isotope is the mass of an atom of the that isotope relative to the mass of an atom of 12C taken as 12 units exactly.” Know that 1.0 amu is defined as exactly 1/12 the mass of a atom. Carbon-12 has 6 protons and 6 neutrons, therefore 1 proton or 1 neutron = ~1 amu 1 amu = x grams

16 Need masses of each isotopes
Average relative atomic mass: is the weighted average for all of the isotopes of a given element, based on the percent abundance of each Need masses of each isotopes Need abundance (percentage) of each isotope This is the value shown on the periodic table

17 Isotopic Masses example…
Chlorine has two isotopes. These have different masses as they have different amounts of neutrons. Using the 12C isotope as a standard, the relative isotopic masses of these two isotopes are (35Cl) and (37Cl). Naturally occurring chlorine is made of 75.80% of the lighter isotope and 24.20% of the heavier isotope.

18 Isotopic Composition of Some Common Elements
ISOTOPES RELATIVE ISOTOPIC MASS ABUNDANCE (%) Hydrogen 1H 1.008 99.986 2H 2.014 0.014 3H 3.016 0.001 Carbon 12C 12 exactly 98.888 13C 13.003 1.112 14C 14.003 Approx 10-10 Oxygen 16O 15.995 99.76 17O 16.999 0.04 18O 17.999 0.20 Silver 107Ag 106.9 51.8 108Ag 108.9 48.2

19 Mass Spectrometer Relative isotopic masses of elements can be obtained using an instrument called a mass spectrometer. This separates the individual isotopes in a sample of the element and determines the mass of each isotope. The information is presented graphically and is known as a mass spectrum.

20 Schematic Diagram of a Mass Spectrometer

21 Vaporization: sample is heated to gas state
Stages Vaporization: sample is heated to gas state Ionization: turned into ions by blasting electrons to knock out electrons from the atoms, creating positively charged ions Acceleration: increases the speed of particles, using an electric field Deflection: by a magnetic field amount of deflection depends on mass and charge of the ion Detection: measures both mass and relative amounts (abundance) of all the ions present

22 Results show the abundance for each isotope of an element
90.92% is neon-20 0.26% is neon-21 8.82% is neon-22

23 Neon Gas

24 Mass Spectrometer cont…
In a mass spectrum showing the isotopes of an element: The number of peaks indicates the number of isotopes The position of each peak on the horizontal axis indicates the relative isotopic mass The relative heights of the peaks correspond to the relative abundance of the isotopes

25 Calculating Relative Isotopic Mass
To calculate average atomic mass you need to know 3 things: # of stable isotopes Mass of each isotope % abundance of each isotope Each isotope is a piece of fruit and the isotope’s mass is the weight of each piece of fruit.

26 Example: Chlorine Calculation
mass of isotope X relative abundance + mass of isotope X relative abundance =_______amu (34.969)(.7553) + (36.935)(.2447) = That’s the same value on the periodic table! Isotope Mass of Isotope Relative Abundance Atomic Mass Cl-35 34.969 75.77% Cl-37 36.935 24.23% amu

27 Example: Copper Calculation
Isotope Mass of Isotope Relative Abundance Atomic Mass Cu-63 amu 69.09% Cu-65 30.91% ( )(.6909)+( )(.3091)= amu

28 Average Relative Mass example…
Imagine taking 100 atoms from a sample of chlorine of chlorine – there will be atoms of 35Cl and of 37Cl. Find the relative atomic mass… Equation to use: ((relative isotopic mass1 x % abundance1) + (relative isotopic mass2 x % abundance2)) /100 OR Ar = ∑(relative isotopic mass x %abundance) / 100

29 Average Atomic Mass cont…
Imagine taking 100 atoms from a sample of chlorine of chlorine – there will be atoms of 35Cl and of 37Cl. Find the relative atomic mass… Ar = Ar =

30 Calculating Relative Abundance
To Calculate % Abundance: Make a Chart Isotopic Mass X %Abundance of each isotope Set-up equation Solve for “x” Plug in “x” value to solve for “y”

31 Example 1.00 10.103 (x) + 11.009 (1 –x) = 10.811 x + y = 1.00
Isotope Mass of Isotope Relative Abundance Atomic Mass B-10 10.013 B-11 11.009 x 1- x 1.00 x + y = 1.00 y = 1 – x (x) (1 –x) = 10.103x x = -0.996x = x = .1987 y= y= .8013 B-10 = 19.87% B-11 = 80.13%

32 Percentage Abundance example…
Copper has two isotopes. 63Cu has a relative isotopic mass of and 65Cu has a relative isotopic mass of The relative atomic mass of copper is Calculate the percentage abundance of the two isotopes. Let x be the percentage abundance of 63Cu So, 100-x is the percentage abundance of 65Cu

33 Percentage Abundance example…
Ar(Cu) = ∑(relative isotopic mass x %abundance) 100 So = 6354 = 62.95x – 64.95x 6354 = 6495 – 2x 2x = 6495 – 6354 2x = 141 x = 70.5

Download ppt "Fundamental Chemical Laws"

Similar presentations

4.1 Defining the Atom Early Models Democritus (about 400BC)- matter composed of tiny, indivisible particles Dalton (about 1800)- - used scientific method.

4.1 Defining the Atom Early Models Democritus (about 400BC)- matter composed of tiny, indivisible particles Dalton (about 1800)- - used scientific method.
Early Models of the Atom –An atom is the smallest particle of an element that retains its identity in a chemical reaction. –Philosophers and scientists.

Early Models of the Atom –An atom is the smallest particle of an element that retains its identity in a chemical reaction. –Philosophers and scientists.
Topic 2.1: Atomic Structure Honors Chemistry 2014-15 Mrs. Peters 1.

Topic 2.1: Atomic Structure Honors Chemistry Mrs. Peters 1.
Atomic Theory of Matter We now take for granted the idea that all matter is comprised of atoms. But how did the Atomic Theory of Matter develop, and how.

Atomic Theory of Matter We now take for granted the idea that all matter is comprised of atoms. But how did the Atomic Theory of Matter develop, and how.
 You cannot see them, yet they make up everything…

 You cannot see them, yet they make up everything…
» What is the name of this device? » Which scientist used this device?

» What is the name of this device? » Which scientist used this device?
Ch. 5: Atomic Structure Standards: Matter consists of atoms that have internal structures that dictate their chemical and physical behavior. Targets:

Ch. 5: Atomic Structure Standards: Matter consists of atoms that have internal structures that dictate their chemical and physical behavior. Targets:
Law of Conservation of Mass:

Law of Conservation of Mass:
Atoms and Atomic Theory Chapter 2. 2.1 Early Chemical Discoveries and the Atomic Theory.

Atoms and Atomic Theory Chapter Early Chemical Discoveries and the Atomic Theory.
Chapter Two: Atoms, Molecules and Ions

Chapter Two: Atoms, Molecules and Ions
Chapter 4: Relative Atomic Mass and the Mole. Masses of Particles Chemists as early as John Dalton, two centuries ago, used experimental data to determine.

Chapter 4: Relative Atomic Mass and the Mole. Masses of Particles Chemists as early as John Dalton, two centuries ago, used experimental data to determine.
Chapter 3 Atoms. People have been thinking about the nature of matter for a long time. The ancient Greeks thought about matter and it wasn’t until the.

Chapter 3 Atoms. People have been thinking about the nature of matter for a long time. The ancient Greeks thought about matter and it wasn’t until the.
History of the Atom; Modern Atomic Theory, Subatomic Particles

History of the Atom; Modern Atomic Theory, Subatomic Particles
THE BUILDING BLOCKS OF MATTER

THE BUILDING BLOCKS OF MATTER
Dr. Namphol Sinkaset Chem 200: General Chemistry I

Dr. Namphol Sinkaset Chem 200: General Chemistry I
New Way Chemistry for Hong Kong A-Level Book 11 Atomic Structure and Relative Masses 1.1The Atomic Nature of Matter 1.2The Experimental Evidence of Atomic.

New Way Chemistry for Hong Kong A-Level Book 11 Atomic Structure and Relative Masses 1.1The Atomic Nature of Matter 1.2The Experimental Evidence of Atomic.
Chapter 2 Atoms, Molecules, Ions.

Chapter 2 Atoms, Molecules, Ions.
Atomic Discovery Early Models of the Atom 400 B.C. – Democritus proposed the existence of fundamental particles of matter that were indivisible and indestructible.

Atomic Discovery Early Models of the Atom 400 B.C. – Democritus proposed the existence of fundamental particles of matter that were indivisible and indestructible.
Atomic Theory Chapter 3 Sections 1 &2 9/18/14.

Atomic Theory Chapter 3 Sections 1 &2 9/18/14.
Chapter 4 Structure of the Atom. History In the 1800’s, early philosophers believed all matter consisted of either air, earth, water, or fire. In the.

Chapter 4 Structure of the Atom. History In the 1800’s, early philosophers believed all matter consisted of either air, earth, water, or fire. In the.

Similar presentations

Ads by Google