Atoms molecules and ions

Chapter 2

Conservation of mass and the Law of definite proportions

• Democritus (460-370 BC) – proposed that elements are composed of tiny particles , atoms from Greek atomos meaning indivisible.

• Robert Boyle (1627 – 1691)- first to study chemistry as a separate discipline and carry out rigorous chemical experiments.

• Joseph Priestly (1733)- isolated oxygen and Antoine Lavoiser(1743 – 1794)- oxygen key substance in combustion, matter is neither created nor destroyed in a chemical reaction.

• Joseph Proust (1754 – 1826)- different samples of a pure chemical substance always contain the same proportion of elements by mass, elements combine in specific proportions, not just random proportions.

Dalton’s atomic theory and the law of multiple proportions

• John Dalton (1766 – 1844)- new theory of matter.1. Elements are made up of tiny particles called atoms. 2. Each element is characterized by the mass of its

atoms. Atoms of the same elements have the same mass, but atoms of different elements have different masses.

3. Chemical combination of elements to make different substances occurs when atoms join together in small whole-number rations. (not fractions)

4. Chemical reactions only rearrange the way atoms are combined, the atoms themselves don’t change.

Law of multiple proportions

• A theory must not only explain observations but predict not yet know. Dalton’s predicts LoMP.

• Elements can combine in different ways to form different substances, whose mass rations are small whole- number multiples of each other.

Practice

• Methane and propane are both constituents of natural gas. A sample of methane contains 5.70 g of carbon atoms and 1.90 g of hydrogen atoms combined in a certain way, whereas a sample of propane contains 4.47 g of carbon atoms and 0.993 g of hydrogen atoms combined in a different way. Show that the two substances obey the law of multiple proportions.

The structure of the atom: electrons

• J.J. Thomson (1856 – 1940) experimented with cathode – ray tubes (sealed vacuum tube with electrode covered with ZnS in it). Cathode ray deflected by a magnet or electrically charged plate. Beam is produced at negative electrode and is deflected to positive plate, must be tiny, negative particles (electrons)

3 factors for electric field

• 1. the strength of the deflecting magnet or electric field. (stronger – greater deflection)

• 2. The size of the negative charge on the electron. (larger charge – greater deflection)

• 3. The mass of the electron. (lighter particle – greater deflection)

Charge to mass ration

• e/m = 1.758820 X 108 C/g

e is the magnitude of the charge in coulomb C, and m is mass in grams.

R. A. Millikan (1868 – 1953) devised a method for measuring the mass of the electron. Oil mist spayed into a chamber allowed to drop between two plates. Mass from falling rate. X-rays to oil made negative and voltage to plates – oil suspended. e was calculated and put above. 9.109 382 x 10 -28 g

The structure of atoms: protons and neutrons

• Ernest Rutherford (1871 – 1937) directed alpha particles (radiation – positive 2x’s electron) at a thin gold foil. He found most passed through (lost of space). Few were deflected (even backwards). The mass in the tiny nucleus. Atom’s diameter 10-10m and nucleus 10 -15m (pea in center of domed stadium) Protons mass – 1.672622 x 10-24 g, Neutron mass – 1.674927 x 10-24

Practice calculations using atomic size

• Ordinary “lead” pencils actually are made of a form of carbon called graphite. If a pencil line is 0.35 mm wide and a diameter of a carbon atom is 1.5 x 10-10 m, how many atoms wide is the line?

Atomic number (z)

• Elements differ according to the number of protons in their atoms’ nuclei.

• Number of protons = number of electrons around atom’s nucleus.

• Mass number (A) = z + N (neutrons)• Except for Hydrogen, most atoms contain as

many neutrons as protons• Isotope – atoms with same number of

protons, but different number of neutron.

Isotope notation• 12

• 6 C

Top number is the mass number, bottom number the atomic number.

Say carbon-12

Isotope practice

• Practice – the isotope of uranium used to generate nuclear power is 235 U. How many protons, neutrons, electrons.

• Element x is toxic to humans in high concentration but is essential to life at low concentrations. Identify element X, whose atoms contain 24 protons, and write the symbol for the isotope with 28 neutrons.

Atomic mass

• Atomic mass unit (amu) aka Dalton (Da) in biological work. 1/12 the mass of C 12 and is equal to 1.600539 x 10-24 g. Protons and neutrons are 1 amu.

• Weighted Average of all the isotopic masses of the element’s naturally occurring isotopes.

• C12 (98.89 % ) and C13 (1.11%). Carbon mass 12.011

Calculating an atomic mass

• Chlorine has two naturally occuring isotope: 35/17 Cl with a natural abundance of 75.77% and an isotopic mass of 34.969 amu and 37/17 with a natural abundance of 24.33% with a mass of 36.966 amu. What is the atomic mass of chlorine?

Compounds

Chemical compound – when 2 or more different elements combine in a specific way to create a new material with different properties than elements alone. Na (soft, silver) Cl ( green gas) NaCl – table salt. Done by chemical reaction. Formula – list symbol each element in there. Subscript – tells the number of each element.

Mixtures

• Simply blends of two or more substances added together in arbitrary proportions, without chemically changing the individual substances. Be separated by physical means.

• Heterogenous – not uniform mixture, salad, muddy water, separates when still

• Homogenous – uniform, constant composition. Air, salt water, koolaid

Covalent bonds: molecules

• 2 elements share (usually 2) electrons. Like a tug-of-war.

• Molecules form this way like: HCl, H2O, NH3

(ammonia), CO2.

• Model using ball and stick (specific covalent bond) or space filling (overall molecular shape)

• Structural formula show specific connection between atoms. H-H

formulas

• Structural formulas useful in organic chemistry.

• Some elements (right corner of periodic table) exists as molecule (H2, N2, O2, F2, Cl2, Br2, I2) are diatomic molecules.

• Practice: Propane has a structure in which the 3 C atoms are bonded and each end C is bonded to 3 H and center C is to 2 H. Draw the structural formula.

Ionic bonds

• Transfer of 1 or more electrons to another usually between metal and nonmetal.

• When atoms lose an electron, it becomes more positive by 1, called cations, metals

• When atoms gain an electron, it becomes more negative by 1, called anions, nonmetals

• Practice: Show the formation of NaCl.

ions

• Can’t really talk about discrete Na+Cl- molecules, but rather an ionic solid. Table salt crystal – cube shape

• Covalent bonded groups of atoms called polyatomic ions have a charge form ionic bonds. When more than one are in a formula, put () around.

• Practice: which is ionic, which molecular

a) BaF2 2) SF4 3) PH3 4) CH3OH

Acids and bases

• Acid - Hydrogen (H+) cation in water (aq) and the anion nonmental / polyatomic ion; hydrochloric (HCl), Nitric (HNO3), sulfuric (H2SO4), Phosphoric acid (H3PO4); the a

• Base - Hydroxide anion (OH-) in water and the metal cation; sodium hydroxide (NaOH or lye or caustic soda), Potassium hydroxide (KOH caustic potash) and barium hydroxide [Ba (OH)2]

Acid / base practice

• Which of the following compounds are acids, and which are bases, explain.

• A) HF

• B) Ca(OH)2

• C) LiOH

• D) HCN

Naming Binary Ionic compounds

• Name must defines it uniquely but also allow chemists (and computers) to know the chemical structure. Organics more complicated. Use simple for now.

• Postive ion takes the name of the element (metal), negative side (nonmetal) gets the –ide ending.

• Ions formed by main groups: 1 (+1), 2 (+2) 3 (+3), 4 (metals) (+4 or +2), 5 (-3), 6 (-2), 7 (-1) Group 8 = 0

Transitional naming

• Transitional carry multiple charges. Fe+2, Fe+3. distinguish by using roman numerals in (). Iron (II), Iron (III). Older forms use ferrous (lower charge), ferric (higher charge).

• Neutral compound: total number of positive charge = total number of negative. One can figure out number of +, by counting the number of negatives on anions.

Practice

• Give systematic names for the following:

a)BaCl2 b) CrCl3 c) PbS d) Fe2O3

Write the formulas for the following:

a)Magnesium fluoride b) Tin (IV) oxide

c) Iron (III) sulfide

Naming binary molecular compounds

• Assume one to be more cation (farther left) like – 1st word and just name, 2nd to be more anion like. Tell how many there are by the subscript using prefixes. Mono – 1, di – 2, tri – 3, tetra – 4, penta – 5, hexa – 6, hepta – 7, octa – 8, nona – 9, dec – 10.

• If prefix ends in a or o and anion name begins with a vowel, drop a or o to avoid 2 vowels together.

Practice

Give systematic names for the following:

• a) PCl3• b) N2O3

• c) P4O7

• d) BrF3

Write formulas for:

a)Carbon dioxide

b)Tetraphoshorus hexachloride

Naming compounds with polyatomic ions

• Named same way as binary ionic. Ba(NO3)2 is called barium nitrate.

• Most have ite or ate ending. (3 , –ide), several form oxoanions (with oxygen). If 2 forms then more O2 gets -ate, less –ite. More than 2 O2 hypo – for least, then –ite, -ate, per – for most

• Will be given a list to memorize.

Practice with polyatomic

• Give systematic names for the following:

a)LiNO3

b)KHSO4

c)CuCO3

d)Fe(ClO4)3

Write the formulas

a)Potassium hypochlorite

b)Silver (I) chromate

c)Iron (III) carbonate

Naming acids

• Most are oxoacid – contain oxygen in addition to hydrogen. –ite gets ous acid and –ate gets ic acid.

• Binary acid use hydro anion ic acid. HCl – hydrochloric acid

• Practice: name the following acids.

a)HBrO (aq)

b)HCN (aq)