Atoms, Molecules, and Ions 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 each element 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) At the same temperature and pressure, equal volumes of different gases contain the same number of particles. 5 liters of oxygen 5 liters of nitrogen 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.

Thomson’s Experiment Voltage source - +

- + Thomson’s Experiment Voltage source Passing an electric current makes a beam appear to move from the negative to the positive end.

Thomson’s Experiment Voltage source By adding an electric field

Thomson’s Experiment Voltage source + - 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 charged particles. 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.

Millikan’s Experiment Atomizer Oil droplets - + Oil Telescope

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 atom Rutherford’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

What he expected

Because

Because, he thought the mass was evenly distributed in the atom.

What he got

How he explained it Atom is mostly empty Small dense, positive particle at center. Alpha particles are 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

A cross section of nuclear atom

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

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 17 17 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 ions Introduction 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.

Molecular formula and structural formula 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

Representing Structure in Molecules Accurately represents the angles at which molecules are attached.

Ions Atoms or groups of atoms with a charge. Cations- positive ions - get by losing electrons(s). Anions- negative ions - get by gaining electron(s). Ionic bonding- Force of attraction between oppositely charged ions. Ionic solids are called salts.

Formation of Cations

Formation of Anions

Examples of ions Cation: A positive ion Mg2+, NH4+ Anion: A negative ion Cl-, SO42- Polyatomic ions

2.7 Introduction to the Periodic Table Elements are classified by: - properties - atomic number Groups (vertical) 1A = alkali metals 2A = alkaline earth metals 7A = halogens 8A = noble gases Periods (horizontal)

Periodic Table Groups /Families Periods

Metals Conductors Lose electrons Malleable and ductile

Nonmetals Brittle Gain electrons Covalent bonds

Semi-metals or Metalloids

Alkali Metals

Alkaline Earth Metals

Halogens

Transition metals

Noble Gases

Inner Transition Metals Lanthanides Lanthanides

Select an element ( ) = Internet link

2.8 Naming Simple Compounds Binary compounds are composed of two electrons Both ionic and covalent compounds will be considered

Binary Ionic Compounds(Type I) 1. Cation first, then anion 2. Monatomic cation takes its name from the name of the element Ca2+ = calcium ion 3. Monatomic anion takes its name from the root of the element + -ide Cl- = chloride CaCl2 = calcium chloride

Name the following compounds CsF Calcium fluoride AlCl3 Aluminum chloride LiH Lithium hydride Calcium hydroxide Ca(OH)2

Some Common Cations

Some Common Anions

Binary Ionic Compounds (Type II): - metal forms more than one cation - use Roman numeral in name PbCl2 Pb2+ is cation PbCl2 = lead (II) chloride PbCl4 = lead (IV) chloride

Example FeCl3 Iron(III) chloride Ferric chloride FeCl2 Iron(II) chloride Ferrous chloride

Example Write the name of the compound Fe2O3 Oxidation state of Fe = +3 Iron(III) oxide Group 1A, Group 2A and Al3+ do not take Roman numerals Silver, Ag forms more than one oxidation state but Roman numerals are not used.

Common Cations and Anions

Ionic compounds with polyatomic ions Polyatomic ions are assigned special names that need to be memorized Oxyanions: anions that contain an atom of a given element and different numbers of oxygen atoms Polyatomic anions (with many atoms) containing oxygen end in -ate or -ite. (The one with more oxygen is called -ate.) Examples: NO3- is nitrate, NO2- is nitrite. (Exceptions: hydroxide (OH-), cyanide (CN-), peroxide (O22-).)

Examples: NaNO3 Sodium nitrate K2SO4 Potassium sulfate Al(HCO3)3 Aluminum bicarbonate or aluminum hydrogen carbonate

Common Polyatomic Ions

Binary Covalent compounds (Type III): Compounds between two nonmetals Although they do not contain ions, they are named similarly to binary ionic compounds - First element in the formula is named first using full element name - Second element is named as if it were an anion. - Use prefixes to denote number of atoms - Never use mono- for naming first element P2O5 = diphosphorus pentoxide

Acids Substances that produce H+ ions when dissolved in water. All acids begin with H. Two types of acids: Oxyacids Non-oxyacids

Naming acids If the formula has oxygen in it write the name of the anion, but change ate to -ic acid ite to -ous acid Watch out for sulfuric and sulfurous HNO3 HNO2 HC2H3O2

Naming acids If the acid doesn’t have oxygen add the prefix hydro- change the suffix -ide to -ic acid HCl H2S HCN