CH.2 Flashcards
Early History of Chemistry
Applications of chemistry before 1000 B.C.
- Use of embalming fluids
- Processing of natural ores to produce metals for ornaments and weapons
The Greeks (400 B.C.)
- Proposed that matter was composed of four fundamental substances (earth, fire, air, and water)
- Considered the question of whether matter is infinitely divisible or is composed of small, indivisible particles
Alchemists dominated the field of chemistry for the next 2000 years
- Helped discover several elements
- Learned to prepare a mineral acids
Modern Chemistry
Foundations were laid in the 16th century by:
- George Bauer, who developed systematic metallurgy
- Paracelsus, who’s discovered the medicinal application of minerals
Robert Boyle (1627-1691)
- Performed quantitative experiments to measure the relationship between pressure and volume of air
- Developed the first experimental definition of an element
- A substance is an element unless it can be broken down into two or more simpler simpler substances
- Held on to certain alchemists’ view
- Metals were not true elements, and eventually, a method to change one metal to another will be found
17th and 18th century
- Interest in the phenomenon of combustion rose
- Joseph Priestly discovered oxygen gas
Georg Stahl
- Suggested that a substance called phlogiston flowed out of burning material
- Postulated that substances that burn in closed containers eventually stop burning since the air in the container is saturated with phlogiston
- Oxygen
Fundamental Chemical Laws
Lavoisier 1771: Law of conservation of mass
Proust 1799: Law of definite proportion
Dalton 1803-1888: Atomic Theory
Law of Conservation of Mass
Total mass of the materials you have before the reaction must be equal the total mass of materials you have at the end
Total mass of reactants = total mass of products
Law of Definite Proportion
Proposed the principle of the constant composition of compounds of Proust’s law ot the law of definite proportions
- Law of definite proportions: a given compound always contain exactly the same proportion of elements by mass
- All samples of a given compound, regardless of their source or how they were prepared, have the same proportions of their constituent elements
Law of Multiple Proportions
When two elements form a series of compounds, the ratios of the masses of the second element that combine with 1 g of the first element can always be reduced to small whole numbers
Dalton’s Atomic Theory
1) Each element is made up of tiny particles called atoms
2) Atoms of a given element are identical
a. Atoms of different elements are different in some fundamental way or ways
3) Chemical compounds are formed when atoms of different elements combine with each other
a. A given compound always has the same relative numbers and types of atoms
4) Chemical reactions involve reorganization of the atoms
a. Atoms themselves are not changed in a chemical reactions
b. Atoms are neither created nor destroyed in chemical reactions
Gay- Lussac and Avogadro (1809-1811)
Joseph Gay-Lussac
- Measured (under the same conditions of T and P) the volumes of gasses that reacted with each other
Avogadro’s Hypothesis
- At the same T and P, equal volumes of different gasses contain the same number of particles
- Makes sense if the distances between the particles in a gas are very great compared with the sizes of the particles
- Volume of a gas is determined by the number of molecules present, not by the size of the individual particles
J.J. Thomson (1900)
Studied electric discharges in partially evacuated tubes called cathode ray tubes
When connected to a high voltage power supply, a glowing area is seen emanating from the cathode
J.J. Thomson cont.
Postulated that the cathode ray was a stream of negatively charged particles (electrons)
- Cathode ray was produced at the negative electrode when high voltage was applied to the tube
- Repelled by the negative pole an applied electric field
- Determined the charge-to-mass ratio of an electron
e/ m = -1/76 x 10^8 C/g
e - charge on the electron (in coulombs)
m - electron mass (in grams)
Thomson’s Results
- The cathode rays are made of tiny particles
These particles have a negative charge
b/c the beam always deflected toward the + plate - The amount of deflection was related to two factors, the charge and mass of the particles
- Every material tested contained these same particles
- The charge/mass of these particles was -1.76 x 10^8 C/g
- The charge/mass of the hydrogen ion is +9.58 x 10^4 C/g
Thomson’s Conclusion
- If the particle has the same amount of charge as a hydrogen ion, then it must have a mass almost 2000x smaller than hydrogen atoms!
- Later experiments by Millikan showed that the particle did have the same amount of charge as the hydrogen ion
- The only way for this to be true is if these particles were pieces of atoms
- Apparently, the atom is not unbreakable
- Thomson believed that these particles were therefore the ultimate building blocks of matter
- These cathode ray particles became known as electrons
(George Stoney, 1874)
Robert Millikan
From 1906-1914 Robert Millikan performed experiments involving charged oil drops, which helped determine the magnitude of electron charge
- He showed ionized oil drops can be balanced against the pull of gravity by an electric field
- The charge is an integral multiple of the electronic charge, e
- Used this value and the charge-to-mass ratio to calculate the mass of an electron is 9.11 x 10^31 kg
Electrons
- Electrons are particles found in all atoms
- Cathode rays are streams of electrons
- The electron has a charge of -1.60 x 10^19 C
- The electron has a mass of 9.1 x 10^-28 g
A New Theory of the Atom
- Since the atom is no longer indivisible, Thomson must propose a new model of the atom to replace the first statement in Dalton’s Atomic Theory
- Rest of Dalton’s theory still valid at this point
Thomson proposes that instead of being a hard, marble-like unbreakable sphere, the way Dalton described it, that it actually had an inner structure
J.J. Thomson: Development of the Plum Pudding Model
- The structure of the atom contains many negatively charged electrons
- These electrons are held in the atom by their attraction for a positively charged electric field within the atom
- There had to be a source of positive charge b/c of the atom is neutral
- Thomson assumed there was no positively charged pieces b/c none showed up in the cathode ray experiment
Predictions of the Plum Pudding Atom
- The mass of the atom is due to the mass of the electrons within it
- The electrons are the only particles in Plum Pudding atoms
- The atoms is mostly empty space
- Cannot have a bunch of negatively charged particles near each other as they would repel
Radioactivity
In the late 1800s, Henri Becquerel and Marie Curie discovered that certain elements would constantly emit small, energetic particles and rays
These energetic particles could penetrate matter
Ernest Rutherford discovered that there were three different kinds of emissions
- alpha, α, particles with a mass 4x H atom and + charge
- beta, β, particles with a mass -1/2000th H atom and - charge
- gamma, γ, rays that are energy rays , not particles
Rutherford’s Experiment
- Carried out to test the accuracy of Thomson’s plum pudding model
- How can you prove something is empty?
- Put something through it
- Involved directing α particles at a thin sheet of metal foil
Expectation
α particles will pass through the foil with minor deflections in their paths
Rutherford’s Experiment: Results
Most α particles, over 98%, passed through the foil
About 2% of the α particles went through but were deflected by large angles
About 0.01% of the α particles bounced off the gold foil
- “…as if you fired a 15” cannon shell at a piece of tissue paper and it came back and hit you
Rutherford’s Experiment Conclusion
Atom mostly empty space
-b/c almost all the particles went straight through
Atom contains a dense particle that was small in volume compared to the atom but large in mass
- b/c of the few particles that bounced back
This dense particle was positively charged
- b/c of the large deflections of some of the particles
Rutherford’s Interpretation: The Nuclear Model
1) The atom contains a tiny dense center called the nucleus
a) The amount of space taken by the nucleus is only about 1/10 trillionth the volume of the atom
2) The nucleus has essentially the entire mass of the atom
a) The electrons weigh so little they give practically no mass to the atom
3) The nucleus is positively charged
a) The amount of positive charge balances the negative charge of the electrons
4) The electrons are dispersed in the empty space of the atom surrounding the nucleus
Structure of the Atom
Rutherford proposed that the nucleus had a particle that had the same amount of charge as an electron but opposite sign
- Based on measurements of the nuclear charge of the elements
These particles are called protons
Charge = +1.60 x 10^19 C
Mass = 1.67262 x 10^-24 g
Since protons and electrons have the same amount of charge, for the atom to be neutral there must be equal numbers of protons and electrons
Radioactive Mass and Charge
It is sometimes easier to compare things to each other rather than to an outside standard
When you do this, the scale of comparison is called a relative scale
We generally talk about the size of charge on atoms by comparing it to the amount of charge on an electron, which we call -1 charge units
- Proton has a charge of +1cu
- Protons and electrons have equal amounts of charge, but opposite signs
We generally talk about the mass of atoms by comparing it to 1/12th the mass of a carbon atom with 6 protons and 6 neutrons, which we call 1 atomic mass unit
- Protons have a mass of 1 amu
- Electrons have a mass of 0.00055 amu, which generally too small to be relevant
There Must Be Something Else There
To answer these questions, Rutherford proposed that there was another particle in the nucleus–it is called a neutron
Neutrons have no charge and a mass of 1 amu
Mass = 1.67493 x 10^-24 g
Slightly heavier than a proton
No charge
Atomic Structure
Nucleus is assumed to contain:
- Protons: have a positive charge that is equal in magnitude to the electron’s negative charge
- Neutrons: have virtually the same mass as a proton but no charge
- Small in size compared to the overall size of the atom
- High density
Atoms of different elements, which have different numbers of protons and electrons, exhibit different chemical behavior
Elements
Each element has a unique number of protons in its nucleus
- The number of protons define the element
The number of protons in the nucleus of an
atom is called the atomic number
- The elements are arranged on the Periodic Table in order of their atomic numbers
Each element has a unique name and symbol
- Symbol either one or two letters
- One capital letter or one capital letter + one lowercase
Isotopes
Atoms with the same number of protons but different numbers of neutrons
- Isotopes of an element have different masses
Depict almost identical chemical properties
In nature, most elements contain mixtures of isotopes
Identifying Isotopes
Atomic number (Z) = number of protons
- Written as a subscript
Mass number (A): total number of protons and neutrons
- Written as a superscript
Abundance = relative amount of found in a sample
Chemical Bonds
Forces that hold atoms together in compounds
Covalent bond: ford by sharing electrons
- Resulting collection of atoms is called a molecule, which can be represented in the following ways:
- Chemical formula
- Structural formula
- Space-filling model
- Ball-and-stick model
Other Methods for Representing a Molecule
Structural formula
- Depicts individual bonds in a molecule
- May or may not indicate the actual shape of the molecule
Space filling model
- illustrates the relative sizes of atom and their relative orientation
Ball-and-stick model
