Multiple Proportions

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Multiple Proportions

• Dalton's law of multiple proportions: When two elements form different compounds, the mass ratio of the elements in one compound is related to the mass ratio in the other by a small whole number.
• Consider the combustion of carbon in the presence of oxygen.
• The carbon is oxidized.
• If there is little oxygen present in the carrier gas the principal product is CO.
  • The mass ratio of O to C in CO is 1.33.
• When oxygen is present in higher concentrations the principal product is CO₂.
  • The mass ratio of O to C in CO₂ is 2.667.
• The ratio of 2.66/1.33 is 2.
  • Ie. the mass ratio of O to C in CO₂ is twice the mass ratio of O to C in CO.

Cathode Ray Tubes

Fig. 2.3 pg. 39

The Discovery of Atomic Structure

• The ancient Greeks were the first to postulate that matter consists of indivisible constituents.
• Later scientists realized that the atom consisted of charged entities.

Cathode Rays and Electrons

• A cathode ray tube (CRT) is a hollow vessel with an electrode at either end.
• A high voltage is applied across the electrodes.
• The voltage causes negative particles to move from the positive electrode to the negative electrode.
• The path of the electrons can be altered by the presence of a magnetic field.

Working of a Cathode Ray Tube

Fig. 2.4 pg. 40

The Discovery of Atomic Structure

Cathode Rays and Electrons

• Consider cathode rays leaving the positive electrode through a small hole.
• If they interact with a magnetic field perpendicular to an applied electric field, the cathode rays can be deflected by different amounts.
• The amount of deflection of the cathode rays depends on the applied magnetic and electric fields.
• In turn, the amount of deflection also depends on the charge to mass ratio of the electron.
• In 1897, Thomson determined the charge to mass ratio of an electron to be 1.76 10⁸ C/g.
• Goal: find the charge on the electron to determine its mass.

Millikan's Oil Drop Experiment

Fig. 2.5 pg. 40

The Discovery of Atomic Structure

Cathode Rays and Electrons

Consider the following experiment:

• oil drops are sprayed above a positively charged plate containing a small hole.
• As the oil drops fall through the hole, they are given a negative charge.
• Gravity forces the drops downward. The applied electric field forces the drops upward.
• When a drop is perfectly balanced, the weight of the drop is equal to the electrostatic force of attraction between the drop and the positive plate.

Millikan carried out the above experiment and determined the charges on the oil drops to be multiples of 1.60 x 10^-19 C.

He concluded the charge on the electron must be 1.60 x 10^-19 C.

Knowing the charge to mass ratio of the electron, we can calculate the mass of the electron to be 9.10939 x 10^-28 g.

Millikan's Oil Drop Experiment (Animation)

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Millikan Oil Drop Experiment

• Robert Millikan (University of Chicago) measured the charge on the electron.
• A fine spray of oil falls through a hole into a chamber where the drops can be observed.
• The plates at the top and bottom of the chamber are charged (the top plate is positive).
• X-rays are shot onto the oil drops which causes the drops to be negatively charged.
• In the absence of voltage, the force on the drops is determined by their mass.
• When a voltage is applied, negatively charged drops will slow down, stop or begin moving upwards. The behavior of the drop is determined by the applied voltage and the charge on the oil drop.
• Millikan used these measurements to determine that the charges on the drops were multiples of 1.6 x 10^-19 C.

He deduced 1.6 x 10^-19 C to be the charge on a single electron.

Thompson's Atomic Model

Fig. 2.9 pg. 42

The Discovery of Atomic Structure

The Nuclear Atom

• From the separation of radiation we conclude that the atom consists of neutral, positively, and negatively charged entities.
• Thomson assumed all these charged species were found in a sphere.
• Therefore, the Thomson model pictures the atom as a sphere with small electrons embedded in a positively charged mass.

Rutherford's Experiment

Fig. 2.10 pg. 42

The Discovery of Atomic Structure

The Nuclear Atom

Rutherford carried out the following experiment:

• A source of alpha-particles was placed at the mouth of a circular detector.
• The alpha-particles were shot through a piece of gold foil.
• Most of the alpha-particles went straight through the foil without deflection.
• If the Thomson model of the atom was correct, then Rutherford's result was impossible.

Rutherford modified Thomson's model as follows:

• assume the atom is spherical but the positive charge must be located at the center, with a diffuse negative charge surrounding it.

Rutherford (Animation)

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Rutherford's Experiment: Nuclear Atom

• Rutherford studied the deflection of alpha-particles by gold foil.
• He found that most of the alpha-particles went through the foil undeflected.
• However, some of the alpha-particles were deflected by large angles (even coming straight back at the source.
• This meant that the alpha-particles had encountered a mass much larger than the He nucleus (alpha-particles are He nuclei).
• However, the mass must have a very small size as only a few alpha-particles encountered them.
• He concluded that the majority of the atom must consist of void space with a small, dense, massive nucleus at the center.

Rutherford Model of the Atom

Fig. 2.11 pg. 43

The Discovery of Atomic Structure

The Nuclear Atom

• In order to get the majority of -particles through a piece of foil to be undeflected, the majority of the atom must consist of a low mass, diffuse negative charge--the electron.
• To account for the small number of high deflections of the -particles, the center or nucleus of the atom must consist of a dense positive charge.

Cross Section of an Atom

Fig. 2.12 pg. 44

The Modern View of Atomic Structure

• The atom consists of positive, negative, and neutral entities (protons, electrons, and neutrons).
• Protons and neutrons are located in the nucleus of the atom, which is small. Most of the mass of the atom is due to the nucleus.
  • There can be a variable number of neutrons for the same number of protons. Isotopes have the same number of protons but different numbers of neutrons.
• Electrons are located outside of the nucleus. Most of the volume of the atom is due to electrons.

Some of the Isotopes of Carbon

Table 2.2 pg. 46

The Modern View of Atomic Structure

Isotopes, Atomic Numbers, and Mass Numbers

• Atomic number (Z) = number of protons in the nucleus.
• Mass number (A) = total number of nucleons in the nucleus (i.e., protons and neutrons).
• By convention, for element X, we write .
• Isotopes have the same Z but different A.

Isotope Table Problem

Problem 2.47 pg. 73

Fill in the gaps in the table shown for problem 2.47 on pg.73.

Symbol	59Co3+	.	.	.
Protons	.	34	76	80
Neutrons	.	46	116	120
Electrons	.	36	.	78
Mass No.	.	.	2+	.

The Mass Spectrometer

Fig. 2.13 pg. 48

The Mass Spectrometer

• Mass spectrometers are pieces of equipment designed to measure atomic and molecular masses accurately.
• The sample is charged as soon as it enters the spectrometer.
• The charged sample is accelerated using an applied voltage.
• The ions are then passed into an evacuated tube and through a magnetic field.
• The magnetic field causes the ions to be deflected by different amounts depending on mass.
• The ions are then detected.

Weighted Averages

Text slide.

Weighted Averages

• To calculate a weighted average of series of values first multiply each value by fraction of total population of objects (ie. atoms/molecules) which give this value when objects are measured.
• (fraction of population measuring value) x (value) = (weighted value)
• Now add up all weighted values to calculate weighted average
• (weighted val.) + (weighted val.) + ... = (weighted avg.)

isotope	24Mg	25Mg	26Mg
abundance	78.70%	10.13%	11.17%
fraction of tot.	0.7870	0.1013	0.1117
mass	23.98504 amu	24.98584 amu	25.98259 amu
weighted mass	18.876 amu	2.5311 amu	2.9023 amu

Periodic Table

Fig. 2.16 pg. 50

The Periodic Table

• The Periodic Table is used to organize the 112 elements in a meaningful way.
• As a consequence of this organization, there are periodic properties associated with the periodic table.
• Columns in the periodic table are called groups (numbered from 1A to 8A or 1 to 18).
• Rows in the periodic table are called periods.
• Metals are located on the left hand side of the periodic table (most of the elements are metals).
• Non-metals are located in the top right hand side of the periodic table.
• Elements with properties similar to both metals and non-metals are called metalloids and are located at the interface between the metals and non-metals.

Groups of the Periodic Table

Table 2.3 pg. 51

The Periodic Table

• Some of the groups in the periodic table are given special names.
• These names indicate the similarities between group members.

Molecular and Empirical Formulas

Text slide.

Molecular and Empirical Formulas

• Empirical formulas experimentally determined by combustion analysis. They give relative numbers of different kinds of atoms in molecule but not necessarily total numbers.

• Molecular formulas give total numbers of different kinds of atoms in molecules. Molecular formulas are exact multiples of empirical formulas.
  

Representation of Molecules

Fig. 2.20 pg. 53

Molecules and Molecular Compounds

Molecules and Chemical Formulas

• A molecule consists of two or more atoms bound together.
• Each molecule has a chemical formula.
• The chemical formula indicates
  1. which atoms are found in the molecule, and
  2. in what proportion they are found.
• Compounds formed from molecules are molecular compounds.

Molecular and Empirical Formulas

• Molecular formulas
  • give the actual numbers and types of atoms in a molecule.
  • Examples: H₂O, CO₂, CO, CH₄, H₂O₂, O₂, O₃, and C₂H₄.
• Empirical formulas
  • give the relative numbers and types of atoms in a molecule.
  • That is, they give the lowest whole number ratio of atoms in a molecule.
  • Examples: H₂O, CO₂, CO, CH₄, HO, O, O, CH₂.
    

Ways of Visualizing Molecules

Fig. 2.21 pg. 54

Molecules and Molecular Compounds

Picturing Molecules

• Molecules occupy three dimensional space.
• However, we often represent them in two dimensions.
• The structural formula gives the connectivity between individual atoms in the molecule.
• The structural formula may or may not be used to show the three dimensional shape of the molecule.
• If the structural formula does show the shape of the molecule, then either a perspective drawing, ball-and-stick model, or space-filling model is used.

Charges on Simple Ions

Fig. 2.22 pg. 57

Ions and Ionic Compounds

• When an atom or molecule loses electrons, it becomes positively charged.
• When an atom or molecule gains electrons, it becomes negatively charged.
• Positively charged ions are called cations.
• Negatively charged ions are called anions.
• An atom or molecule can lose more than one electron.
• The number of electrons an atom loses is related to its position on the periodic table.
• Metals tend to form cations whereas non-metals tend to form anions.

Formation of NaCl

Fig. 2.23 pg. 57

Ions and Ionic Compounds

Ionic Compounds

• The majority of chemistry involves the transfer of electrons between species.
• Example:
  • To form NaCl, the neutral sodium atom, Na, must lose an electron to become a cation: Na⁺.
  • The electron cannot be lost entirely, so it is transferred to a chlorine atom, Cl, which then becomes an anion: Cl^-.
  • The Na⁺ and Cl^- ions are attracted to form an ionic NaCl lattice which crystallizes.
• Important: note that there are no easily identified NaCl molecules in the ionic lattice. Therefore, we cannot use molecular formulas to describe ionic substances.

Consider the formation of Mg₃N₂:

• Mg loses two electrons to become Mg²⁺
• Nitrogen gains three electrons to become N^3-.
• For a neutral species, the # of electrons lost and gained must be equal.
• However, Mg can only lose electrons in twos and N can only accept electrons in threes.
• Therefore, Mg needs to lose 6 electrons (2x3) and N gain those 6 electrons (3x2).

• Therefore, the formula is Mg₃N₂.

Na + Chlorine Reaction

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Formation of NaCl

• Sodium is melted in a crucible and placed in green chlorine gas.
• The reaction is highly exothermic:

2Na(l) + Cl₂(g) 2NaCl(s),

• Chlorine reacts with sodium producing a bright yellow flame.
• At the end of the reaction, only white sodium chloride remains.

Common Cations

Table 2.4 pg. 62

Naming Inorganic Compounds

• Naming of compounds, nomenclature, is divided into organic compounds (those containing C) and inorganic compounds (the rest of the periodic table).

Names and Formulas for Cations

1. Cations formed from a metal have the same name as the metal.

2. If the metal can form more than one cation, then the charge is indicated in parentheses in the name.

3. Cations formed from non-metals end in -ium.

Naming Anions

Fig. 2.26 pg. 62

Naming Inorganic Compounds

Names and Formulas for Anions

• Polyatomic anions containing oxygen with more than two members in the series are named as follows (in order of decreasing oxygen):
  1. per-?-ate
  2. -ate
  3. -ite
  4. hypo-?-ite
• Simple anions are given the suffix -ide.

Ionic Compounds

• These are named cation then anion.
• Example: BaBr₂ = barium bromide.

Common Anions

Table 2.5 pg. 64

Naming Inorganic Compounds

Names and Formulas for Anions

1. Monatomic anions (with only one atom) are called -ide.
   • Example: Cl^- is chloride ion.
   • Exceptions: hydroxide, cyanide and peroxide ions.
2. 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, SO₄^2- is sulfate,SO₃^2- is sulfite
3. Polyatomic anions containing oxygen with more than two members in the series are named as follows (in order of decreasing oxygen):
   • per-?-ate
   • -ate
   • -ite
   • hypo-?-ite
4. Polyatomic anions containing oxygen with additional hydrogens are named by adding hydrogen or bi- (one H), dihydrogen (two H), etc., to the name as follows:
   • CO₃^2- is the carbonate anion
   • HCO₃^- is the hydrogen carbonate (or bicarbonate) anion.
   • H₂PO₄^- is the dihydrogen phosphate anion.
   • HPO₄^2- is the hydrogen phosphate anion.

Naming Acids From Anions

Fig. 2.20 pg. 65

Naming Inorganic Compounds

Names and Formulas of Acids

The names of acids are related to the names of anions:

• -ide becomes hydro-?-ic acid;
• -ate becomes -ic acid;
• -ite becomes -ous acid.

Naming Inorganic Compounds

Text slide.

Naming Inorganic Compounds

Ionic Compounds

• Composed of metal or polyatomic cations plus nonmetal or polyatomic anions
• Metal named before nonmetal
• If metal can have multiple oxidation numbers use Roman numerals in parentheses

Covalent Binary Compounds

• Composed of two kinds of nonmetallic atoms
• Prefix both atom names with Latin number prefixes = molecular formula subscripts (mon, di, tri, etc.)
• Atoms leftmost in periodic table named first

Naming Binary Compounds

Table 2.6 pg. 66

Naming Inorganic Compounds

Names and Formulas for Binary Molecular Compounds

• Binary molecular compounds have two elements.
• The most metallic element is usually written first (i.e., the one to the farthest left on the periodic table). Exception: NH₃.
• If both elements are in the same group, the lower one is written first.
• Greek prefixes are used to indicate the number of atoms.

Examples:

Cl₂O is dichlorine monoxide,

N₂O₄ is dinitrogen tetroxide,

NF₃ is nitrogen trifluoride,

P₄S₁₀ is tetraphosphorus decasulfide.