Chapter 2 Atoms, Molecules, and Ions
Chapter 2: Topics • Early history of chemistry • Fundamental chemical laws • Dalton’s atomic theory • Early experiments to characterize the atom • The modern view of atomic structure • Molecules and ions • An introduction to the Periodic Table • Naming simple compounds
2.1 The early history of chemistry • Greeks • Democritus and others - atomos • Alchemy • 1660 - Robert Boyle- experimental definition of element. • Lavoisier- Father of modern chemistry.
Greeks • Matter is composed of fire, earth, water and air • The Greek philosopher Democritus (460 B.C. – 370 B.C.) was among the first to suggest the existence of atoms (from the Greek word “atomos”) • He believed that atoms were indivisible and indestructible • His ideas did agree with later scientific theory, but did not explain chemical behavior, and was not based on the scientific method– but just philosophy
Alchemy • Turning Cheep metals into gold • Alchemists discovered several elements and prepared mineral acids
17th Century • Robert Boyle: First “chemist” to perform quantitative experiments • He published his first book: “The Skeptical Chemist” in 1661. • He talked about elements
18th Century • George Stahl: Phlogiston flows out of a burning material. • Joseph Priestley: Discovers oxygen gas, “dephlogisticated air, i.e., low in phlogistone”
2.2 Fundamentals chemical Laws • Law of Conservation of Mass • Law of Definite Proportion • Law of Multiple
Law of Conservation of Mass • It was discovered by Antoine Lavoisier • It was the basis for development of chemistry in the 19th century • Mass is neither created nor destroyed • Combustion involves oxygen, not phlogiston
Law of Definite Proportion(Proust’s Law) • A given compound always contains exactly the same proportion of elements by mass. • Water is composed of 11.1% H and 88.9% O (w/w)
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 gram of the first element can always be reduced to small whole numbers. • The ratio of the masses of oxygen that combine with 1g of H in H2O and H2O2 will be a small whole number (“2”).
Example • Water, H2O has 8 g of oxygen per 1g of hydrogen. • Hydrogen peroxide, H2O2, has 16 g of oxygen per 1g of hydrogen. • 16/8 = 2/1 • Small whole number ratios. • This fact could be explained in terms of atoms
2.3 Dalton’s Atomic Theory (1808) • Elements are made up of small particles called atoms • Atoms of eachelement are identical. Atoms of different elements are different. • Compounds are formed when atoms combine. Each compound has a always same type and relative number of atoms • Chemical reactions are rearrangement of atoms but atoms are never changed into atoms of other element. , or created or destroyed.
Gay-Lussac hypothesis (1809) • Provided basics to determining absolute formulas of compounds • Gay-Lussac- under the same conditions of temperature and pressure, compounds always react in whole number ratios by volume. • 2volumes of H react with one volume of O to form 2volumes of gaseous water and
Avogadro’s Hypothesis (1811) • 5 liters of oxygen • 5 liters of nitrogen • Same number of particles! At the same temperature and pressure, equal volumes of different gases contain the same number of particles.
If Avogadro's hypothesis is correct, Gay-Lussac’s can be interpreted as follows: • 2 molecules of H react with 1 molecule of O 2 molecules of H2O
2.4 Early experiments to characterize the atom • Based on Dalton, Gay-Lussac, Avogadro and others, work started to identify the nature of the atom • What is an atom made of? How do atoms of various elements differ?
The electron • J. J. Thomson - postulated the existence of electrons using cathode ray tubes. • Ernest Rutherford - explained the nuclear atom, containing a dense nucleus with electrons traveling around the nucleus at a large distance.
Thomson’s Experiment Voltage source - + When high voltage is applied to the tube a ray emanates from the cathode is called cathode ray.
Voltage source Thomson’s Experiment - +
Voltage source Thomson’s Experiment - + • Passing an electric current makes a beam appear to move from the negative to the positive end.
Voltage source Thomson’s Experiment • By adding an electric field
Voltage source Thomson’s Experiment + - • By adding an electric field, he found that the moving particles were negatively charged
Results of Thomson Experiment • Electrons are produced from electrodes made from various types of metals, all atoms must contain electrons. • Since atoms are electrically neutral, they must contain positively chargedparticles. • Thomson determined charge-to-mass ratio of an electron: • e/m = -1.76X108C/g
Thomson’s Model • Atom consisted of a diffuse cloud of positive charge with negative electrons embedded randomly • Atom was like plum pudding. • Thomson believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model.
Atomizer Oil droplets + - Oil Telescope Millikan’s Experiment
Millikan’s Experiment X-rays X-rays give some electrons a charge.
Millikan’s Experiment From the mass of the drop and the charge on the plates, the mass of an electron is calculated
Radioactivity • Certain elements produce high energy radiation • Discovered by accident and was a result of spontaneous emission by uranium • Bequerel (1896) found that a piece of mineral containing uranium could produce an image on a photographic plate in the absence of light. • Three types of radiation were known: • alpha- helium nucleus (+2 charge, 7300 times that of the electron) • beta- high speed electron • gamma- high energy light
The nuclear atomRutherford’s Experiment • Aimed at testing Thomson’s plum pudding model • Used uranium to produce alpha particles. • Alpha particles are directed at gold foil through hole in lead block. • Since the mass is evenly distributed in gold atoms alpha particles should go straight through. • Used gold foil because it could be made atoms thin.
Florescent Screen Lead block Uranium Gold Foil
+ How he explained it • Atom is mostly empty • Small dense, positive particle at center. • Alpha particlesare deflected by it if they get close enough.
+ Proof for nuclear atom
Nuclear atom model • According to Rutherford: The atom consists of a dense center of positive charge (Nucleus) with electrons moving around it at distance that is large relative to the nuclear radius
2.5 The modern view of an atomic structure:An introduction • The atom is mostly empty space. • Two regions • Nucleus- protons and neutrons. • It is characterized by small size and high density • Electron cloud- region where you might find an electron. • The chemistry of atom • Results mainly from electrons
Mass and charge of nuclear particles Particle Mass (Kg) Charge Electron 9.11X10-31 -1 Proton 1.67X10-27 +1 Neutron 1.67X10-27 None
Why atoms of different elements have different properties? • Atoms of different elements have different number of protons and electrons. • Number and arrangement of electrons around nucleus differ from one element to another.
Sub-atomic Particles • Z - atomic number = number of protons determines type of atom. • A - mass number = number of protons + neutrons. • Number of protons = number of electrons if atom is neutral.
Symbols A Mass number X Atomic number Z 23 Na 11 Na-23
More Atomic Symbols 16 31 65 O P Zn 8 15 30 8 p+ 15 p+ 30 p+ 8 n 16 n 35 n 8 e- 15e- 30 e-
Isotopes • Atoms of the same element (same atomic number) with different mass numbers • Atoms with the same number of protons, but different numbers of neutrons. Isotopes of chlorine 35Cl 37Cl 1717 chlorine - 35 chlorine – 37 Cl-35 Cl-37
Two isotopes of sodium • Isotopes show almost identical chemical properties. Why? • They possess same number of electrons
2.6 Molecules and ionsIntroduction to chemical bonding • The forces that hold atoms together are called chemical bonding • Covalent bonding - sharing electrons. • Collection of atoms by covalent bonding lead to molecules • Molecules can be represented by formulas • Chemical formula- Symbol relates number and type of atoms in a molecule. • Diatomic molecule: two atoms of same element are connected by a covalent bond.
Molecules that contain two atoms of the same element bonded together are called diatomic molecules.
Molecular Compounds • Molecular formulas • give the actual numbers and types of atoms in a molecule. • Examples: H2O, CO2, CO, CH4, H2O2, O2, O3, and C2H4. • Structural formula: bonds are shown as lines H H H C C H H H