Chapter 2 – Atoms and Elements What is chemistry? “A branch of science which deals with the elementary substances or forms of matter, of which all bodies are composed, the laws that regulate the combination of these elements in the formation of compound bodies, and the phenomena that accompany their exposure to diverse physical conditions”. How does it interact with others or react with its surroundings ? Composition Preparation Reaction What is it made of ? How is it made?
Transcript
Chapter 2 – Atoms and Elements
What is chemistry?
“A branch of science which deals with the elementary substances or forms of matter, of which all bodies are composed, the laws that regulate the combination of these elements in the formation of compound bodies, and the phenomena that accompany their exposure to diverse physical conditions”.
How does it interact with others or react with its surroundings ?
Composition
Preparation
Reaction
What is it made of ?
How is it made?
Ex) COFFEE
Composition:
i) Organic Compounds: ii) Inorganic Compounds:
Proteins
Acids
Esters
Sugars
Caffeine
Pesticides
Dissolved Salts
Dissolved minerals
Water
Preparation
Grown
• Biochemical processes make the organic & biological compounds
Roasted
• Heat combined with air burns off undesired compounds & converts some to those that give flavor
• Caffeine is burned off if roasted too long
• Decaffeination
• Hot water poured over powder, where all water soluble compounds dissolve. The liquid is separated from the bean residue by a filtration.
Preparation
• Pulverization of the bean to increase the surface are to aid extraction process
Ground
Extraction
• Makes it more vulnerable to oxidation affecting taste & shelf life
• When it sits the element exposed to the air the organic compounds oxidizes causing a bitter taste.
Reaction with Surroundings
• Stimulant – increases heart rate by promoting adrenaline production
Caffeine
Burns
• Diuretic – stimulates urine production
Where does chemistry fit?
What is the purpose of modern chemistry?
1) Physical chemistry/chemical physics
Thermodynamics
Kinetics
Spectroscopy
Quantum Mechanics
2) Analytical/Environmental Chemistry
Quantification and identification techniques
Separation Methods
Development of New instrumentation
Forensic/Environmental Chemistry
3) Preparative Chemistry
Synthesis of new organic compounds
Material Science
Pharmaceutical Chemistry
Synthesis of new inorganic compounds
4) Biochemistry/Molecular Biology
Chemical processes of Life
Exploration of DNA
Rational Drug Design
Structure and function of proteins
Atomic Theory
Greeks
Aristotle
Plato
Atom ( A – not, tomos – to cut)
- Revelation of truth through logic
- Cosmic order - Hierarchy of being
Atomic Theory
Greeks
Five perfect shapes
TetrahedronCubeOctahedronDodecahedronIcosahedron
Five elements
FireWaterWindEarthEther Technology
Steam Engines Organs Jewelry Reinforced Concrete
Medieval Times
Alchemy
Transmutation of lead into gold.
“Understanding God’s Creation”“Moral Teachings From Nature”
Religious Society
Communal
Static
God Centered , Hierarchical
No Individual Identity
Cycles of Life and NaturePreordained unchangeable OrderNo Change
Chemical knowledge of the timeCoal/Peat Fuel
Beer/Wine
Fragrances/Extracts
Glass/Ceramics
Metal work: - steel - pewter - quicksilver - jewelry
Poisons/Medicines
Pigments/Dyes
Salt Production
Distillation
Enlightenment
Scientific Method
Determinism
Times of Change/Discovery
“ Mechanistic Understanding of the Universe”
- French and American Revolution.
- Industrial Revolution
Mechanistic Thinking
Materialism Earth Centered
Individualism Career Scientist
- Rapid exploration of chemistry began: New Elements Natural Products Synthetic Methods
Lavoisier 1785 “Conservation of mass”
Joseph Proust 1794 “Law of Definite Proportions”
John Dalton 1808 “Atomic Theory of Matter”
1. All matter consists of solid and indivisible atoms.
2. All of the atoms of a given chemical element are identical in mass and in all other properties.
3. Different elements have different kinds of atoms; these atoms differ in mass from element to element.
4. Atoms are indestructible & retain their identity in all chemical reactions.
5. The formation of a compound from its elements occurs through the combination of atoms of unlike elements in small whole-number ratios.
Modifications Required to Daltons Theory
1. Atoms can be further divided into subatomic particles.
2. Different isotopes of an element have different masses
3. Valid, However some have very similar masses.
4. In nuclear reactions, atoms do not retain their identity.
5. Valid, however, Dalton was unaware that not all elements are made up of single atoms.
Ex) Radium-226 → Radon-222 + a-4
Ex) Protons, neutrons, electrons
Ex) Carbon-12 12.000 u Carbon-13 13.003 u Carbon-14 14.003 u
Ex) Nitrogen-14 14.003 u. Carbon-14 14.003 u.
Ex) Chlorine is a gas composed of diatomic molecules Cl2.
Ex) Bromine is a liquid composed of diatomic molecules Br2.
Elemental Forms
Molecular Gas
Molecular Liquid
H2, N2, O2, F2, and Cl2.
Elemental Forms
Molecular Solids
Atomic Liquids
Ex) I2, P4, and S8.
Atomic Gases
Ex) He, Ne, Ar, Kr, Xe, and Rn .
Ex) Hg, Ga
Infinite Networks
Ex) Metals Diamond Graphite
Graphite (2-D) Diamond (3-D)
Compounds
Molecular
Infinite networks – SiO2 Sand, Glass
Ionic
Complex ions are also possible such as:
SO32-, NO3
-, PO4-
NH4+, H3O+
Properties
Chemical Physical
Is observed without changing a compound/element into another compound/element
Is observed by changing a compound/element into another compound/element.
Melting point
Freezing Point
Density
Chemical Reactions
Energy of Reaction
Combustion
Modern Atomic Theory
In the late 19-th and early 20-th century the basic principles of modern atomic theory were laid down
Radioactivity
Electron
Proton/Nucleus
J.J. Thomson 1896
R. A. Millikan 1909
Henri Becquerel 1896
Marie and Paul Currie 1899
Ernest Rutherford 1919
Neutron J. Chadwick 1932
Electrons
Anode: positive electrode
Cathode Ray Tube
Current flows when tube is evacuated
Hole drilled in tube. Gass entering tube glows
Cathode: negative electrode
Cathode Rays
Electron charge-to-mass ratioJ.J. Thomson – 1897 - cathode rays are negatively
charged particles
CRT with electric and magnetic fields applied at right angles
Beam deflects to positively charged plate
Magnetic field applied to deflected beam
Changes in the deflection behaviour allowed the mass to charge ratio of the electron to be determined at 1.7588202 C/kg
Oil Drop ExperimentR Millikan and H A Fletcher (1909)
Accurate measurement of the electron charge.
Balanced the force of gravity with an opposing electric force
The balancing force between droplets had common factor
He surmised that the charge of a single electrone = 1.60217646 10-19 C
Applying the charge/mass ratio,mass of e = 9.1093819 10-31 kg
“Canal Rays” and Protons
+Anode
-Cathodee-
e-
e-
++
+
E Goldstein (1850-1930) discovered canal rays in 1886using a “reverse cathode ray” tube
Electrons emitted from the cathode hit gas molecules causing ionization into (more) electrons and leaving positively charged “ions” which travel to the cathodeThose that pass through
the hole (“canal”) can be analyzed for charge-mass ratio, which are much smaller than electron, but largest for hydrogen
E. Rutherford determined that the hydrogen cation is a fundamental particle, and named it the proton
Radioactivity
Three fundamental types of nuclear radiation were identified by how they respond to electric fields by E. Rutherford.
Paul and Marie Currie isolated the radioactive elements Radium and Polonium. They postulated that their spontaneously emitted radiation was the result of nuclear disintegration.
Three types of radiation:
alpha, , beta, , and gamma, .
Radioactivity: propertiesFrom their charge-mass ratios and other experiments of these rays were characterized and identifiedAlpha particles: He2+ nuclei m = 4 amu q =+2)
Beta particles: electron (e-) (identical to cathode rays)
Gamma rays: high-energy light, with wavelengths shorter than X-rays
Rutherford experimentUsing alpha particles, he bombarded a very thin foil of gold and observed deflections using a circular fluorescent screen
The nuclear atom
Rutherford said of the alpha particles deflected almost straight back.
He tried to prove the plum pudding model of the atom propose by Thomson, which is composed of electrons imbedded in a sphere of uniform positive charge.
Deflection angle and frequency were carefully measured, which led to the conclusions:
1. Most of gold foil is empty space2. There are small centers of highly-positive charge3. Centers have high mass to resist displacement4. Size of atom estimated from distance between centers to be ~10-10 m diameter.5. Size of centers estimated to be ~10-15 m diameter
Centers were called the nucleus.
Electrons occupy the volume of the atom outside the nucleus
Constituents of the atomIn 1920 Rutherford predicted the existence of the neutralparticle with mass equal to that of a proton and electron.
In 1932 Chadwick verified experimentally the existence of the neutron
Relative mass of carbon defined t be 12 u
The mass spectrometerMass spectrometer is a variation on the CRT, developed by J.J. Thomson, which allows the determination of m/z ratios of cations.
Cations of differing m/z ratio’s can be selected by adjusting the magnetic field strength
Average atomic mass
35Cl has 17 protons and 18 neutrons
37Cl has 17 protons and 20 neutrons
Isotopes are atoms of the same element that differ in mass due to differences in the number of neutrons
The atomic mass of Chlorine is a weighted average between the two isotopes as:
Atomic Mass = Mass(Cl-35) *frac.(Cl-35) + Mass(Cl-37) *frac.(Cl-37) = (34.968)*(0.7537) + (36.956)*(0.2463) = 35.46 u
Defining an Element The atomic mass unit (u) is defined as one twelfth of the mass of a carbon atom containing six protons, six neutrons and six electrons: 1 u = 1.661 × 10-24 g
The mass of an atom in u will be approximately equal to the combined number of protons and neutrons it contains.
C12
6
mass numbersymbol
atomic numberAtomic number (Z) = # protons
Mass number (A) = # protons + # neutrons
The atomic # determines the identity of the element (optional).
If # p’s = #e’s neutral
If # p’s > # e’s cation
If # p’s < # e’s anion
e.g. Gallium has two naturally occurring isotopes and an average atomic mass of 69.723 u:
Calculate the percent abundance of each isotope of gallium.
The MoleIt is not practical to work on the scale of individual atoms. It is necessary to work on the macroscopic scale.
It was found that for 6.0221*1023 atoms for any element the mass corresponds to the atomic mass in grams.
Ex) 12.00 g of carbon corresponds to 6.02221*1023 carbon atoms
The same is true for molecules.
Ex) CO2 weighs 12.000 + 2*15.999 = 43.998 u 6.02221*1023 molecules of CO2 weighs 43.998 g
“1 mole = 6.0221 × 1023”
This number, 6.0221*1023, was named after Amedeo
Avogadro who initialyl proposed the idea.
Molar MassThe molar mass of a particle is the mass in grams of one mole, 6.0221*1023, particles
Ex) CO2 has molecular mass of 43.998 u therefore, it has a molar mass of 43.998 g/mol
Exercise: How many moles of Cl2 are there in 105.7g.
The atomic mass of Cl is 35.46 u
# moles Cl2 = mass Cl2/molar mass Cl2 =105.7g/70.92 g/mol = 1.490 moles
The molecular mass of Cl2 is 2*35.46 u = 70.92 u
Its molar mass is 70.92 g/mol.
Ex ) One mole of protons weighs
(6.0221*1023)*(1.67*10-24 g) = 1.01 g
Exercise: What is the mass of carbon dioxide containing 2.57*1021 atoms of oxygen
CO2 contains two O atoms for every CO2 molecule
# CO2 = (# of O)/2 = 2.57*1021/2 = 1.29*1021
# moles CO2 = (# CO2 molecules )/( 6.0221*1023 molecules/mol)
# moles CO2 = (1.29*1021 molecules)/( 6.0221*1023 molecules/mol)
# moles CO2 = 0.00214 mol
How much does this weigh?
mass CO2 = (# moles CO2)*(molar mass of CO2)
mass CO2 = (0.00214 mol)*(43.998 g/mol) = 0.0916 g
Law of Periodicity “The properties of the elements areperiodic functions of atomic number.”
Metals – Conducting, Ductile
Metalloids - Semiconductors Ductile ?
Nonmetals – insulatorsnot ductile
Group Period
Repetition of properties
Similar chemical properties
Overview of the Elements by GroupHydrogen (H)
Has properties of groups 1 and 17 but doesn’t belong to either.Diatomic gas (H2)
Group 1: Alkali Metals (Li, Na, K, Rb, Cs, Fr)
Soft metals that react strongly with Water and oxygen. (reactivity increases with atomic mass)
Do not exist in pure form in nature dueto high reactivity
Readily lose 1 electron to make cations with +1 charge
Unreactive
Group 2: Alkaline Earth Metals (Be, Mg, Ca, Sr, Ba, Ra )
Most are metals that react with water to give X(OH)2 and with oxygen to give XO.
Reactivity increases withatomic mass.
Beryllium does not react with water (Highly toxic)
Do not exist in pure form in nature due to high reactivity
Readily lose 2 electrons to make cations with +2 charge
Groups 3-12: Transition Metals
Group 3 metals lose 3 e- to make +3 cations otherwise act like Group 2 metals
Groups 4-11 are the ‘true’ transition metals in that they lose e- to form coloured compounds in which the metal atom has a positive charge
Groups 11 and 12 mimic the behaviour of Groups 1 and 2 respectively but are less reactive
Copper and gold are the only coloured metals
Silver is the best conductor
Gold is the most malleable metal
Readily lose electrons to make cations
Mercury is the only liquid metal;the rest are solids
The metals in the middle of thetransition groups are hardest Ex) group 6 Cr, Mo, W
The metals at the edges arethe softest (Groups 3 and 12)Ex) Zn and Ga
Group 13 (B, Al, Ga, In, Tl)
Most are metals; Boron is a metalloid
All are solids, Gallium -low m.pt.
Aluminum - industrially important- third most abundant- produces its own protective layer
Lose 3 electrons to make cations with +3 charge
Form compounds in a 1:3 ratio with halogens (e.g. BCl3)
C is a nonmetal; Si and Ge are metalloids;Sn and Pb are metals
Si is the 2-nd most abundant Element which does occur in pure form but as silicates(compounds made of siliconand oxygen) which form rocks,sand, glass, etc.
Form compounds in a 1:4 ratio with halogens (e.g. CCl4)
Group 14 (C, Si, Ge, Sn, Pb)
Carbon exists in several different Allotropes: 1) graphite2) diamonds3) fullerenes – Many types
Carbon is the backbone atom of organic and biological molecules
CARBON
Group 15: Pnictogens (N, P, As, Sb, Bi)
N and P are nonmetalsAs and Sb are metalloidsBi is somewhat metallic
N is a highly stable diatomic gas (N2)and the most abundant element in the atmosphere
P in three allotropes White, P4 – Fire BombsRed and black, polymers - used in match heads)
Form compounds in a 1:3 ratio with hydrogen (e.g. NH3)
Group 16: Chalcogens (O, S, Se, Te, Po)O, S & Se are nonmetalsTe is a metalloid Po is a metal
O is the most abundant in the earth’s crust & the second most in the atm.
O exists in two allotropes : O2 and O3 both are very reactive gases
S exists in many allotropes: S2, S6, S8, etc.
Form compounds in a 1:2 ratio with hydrogen (e.g. H2O)
Gain 2 electrons to make anions with -2 charge
Group 17: Halogens (F, Cl, Br, I, At)Nonmetals that exist as diatomic molecules (except for astatine which is too unstable to study)
F & Cl are gasesBr is a liquidIodine is a solid
colourful F2 is yellowCl2 is yellow-greenBr2 is red-brownI2 is dark purple
Form compounds in a 1:1 ratio with hydrogen (e.g. HF)
Gain 1 electron to make anions with -1 charge
F most reactive known
Cl2
Br2l2
Group 18: Noble Gases (He, Ne, Ar, Kr, Xe, Rn)
Inert- unreactive gaseous nonmetals
Exist primarily in elemental from
Compounds have beenmade containing Xe Ex) XeF2
Helium low density – BalloonsLow BP - Coolant
XenonUsed as a probe to study structure in porous material
