Atoms, molecules and ions

1. Atoms, molecules and ions

2. Dalton’s Atomic Theory (1808) • The elements are formed of extremely small particles, called atoms. All the atoms of a given element are identical. The atoms of an elements are different than those of all other elements. • Compounds are formed of more than one element. In all compounds, the ratio between the number of atoms of two elements is either an integer or a simple fraction .

3. Dalton’s Atomic Theory (1808) • A chemical reaction is the separation, combination, or rearrangement of atoms; it causes neither the destruction nor the creation of atoms. • N.B. Dalton did not know the structure of the atom (i.e., electron, proton, neutron, nucleus, etc.) - Dalton imagined an atom was small and indivisible

4. Dalton’s Atomic Theory (1808) • Dalton’s work explained some empirical observations • The law of definite proportions • The law of multiple proportions • The law of conservation of mass

5. The Structure of the Atom: the Electron • With a cathode ray tube, the existence of a subatomic particle with a negative charge was discovered by Thomson…the electron. • Thomson established the charge/mass ratio of the electron.

6. The Structure of the Atom: the Electron • Millikan established the charge of an electron in the early 1900’s (-1.60 x 10-19 C) • With the help of Thomson’s work, Millikan deduced the mass of an electron

7. Radioactivity • Radioactivity: the spontaneous emission of particles and/or radiation • Three types of particle/rays are produced by the decay of radioactive substances : • αparticles(helium nuclei) • βparticles(electrons) • γ rays (rays of high energy, and no charge)

8. The Structure of the Atom: the Proton and the Nucleus • Before 1910, Thompson’s “plum pudding” model of the atom was the accepted model. • In this model, the electrons are dispersed within a much bigger sphere of uniform and positive matter. • This is “logical" considering the small mass of the electron compared to the overall mass of the atom

9. The Structure of the Atom: the Proton and the Nucleus • Rutherford bombarded thin sheets of gold with α particles • According to the “plum pudding” model, α particles should pass through without their trajectories being greatly affected • However, the trajectories of some particles are greatly perturbed

10. The Structure of the Atom: the Proton and the Nucleus • In Rutherford’s experiment, the deviations are the result of large repulsions • Because α particles are positively charged, the positive charge in the atom must be concentrated in a very small and solid nucleus in the center of the atom (about 10-13 of the volume of the atom, but almost 100% of the mass) • The positive charges that are located in the nucleus are the protons • The charge of a proton is the same as the electron, but with the opposite sign (i.e., positive) • The mass of the proton is 1.67252 x 10-24 g (about 1840 times that of the electron)

11. The Structure of the Atom: the Neutron • Before 1932, a great mystery was the fact that H had one proton and He had two, but He was four times more massive than H! • Many people predicted the existence of the neutron, but it was Chadwick that discovered it in 1932 • The neutron has no charge and is slightly larger than the proton (1.67493 x 10-24 g for the neutron vs 1.67252 x 10-24 g for the proton)

12. Atomic number, Atomic Mass, and Isotopes • Atomic number: The number of protons in the nucleus of each atom of a given element • Atomic mass: The total number of neutrons and protons contained in the nucleus of an atom • All atoms of an element have the same number of protons, but not necessarily the same number of neutrons • Atoms that have the same atomic number but different atomic masses are called isotopes

13. Atomic number, Atomic Mass, and Isotopes • To identify an isotope, we use the notation , where A is the atomic mass, Z is the atomic number and X is the element in question • Isotopes of an element have essentially the same chemistry since the number of neutrons has little influence on the chemistry of an element • We use the atomic mass to identify isotopes (e.g..; is uranium 235 and is uranium 238) • Hydrogen has three isotopes (only hydrogen has special names for its isotopes) hydrogen: deuterium: tritium:

14. The Periodic Table • Table in which elements are grouped based on similar physical and chemical properties. • The horizontal rows are called periods • The vertical columns are called groups (or families) • Three categories of elements: • metal (good conductor) • non-metal (poor conductor) • metalloid (between metal and non-metal)

15. Molecules and Ions • A molecule is an assembly of at least two atoms held together in a fixed arrangement, by chemical forces • An ion is an atom or molecule that has gained or lost electrons • Cation: ion with a positive charge • Anion: ion with a negative charge • An ionic compound is a composed of cations and anions • e.g..; NaOH(s) is a group of Na+ and OH- ions and there are no distinct molecules of NaOH

16. Chemical Formulas • Molecular formula: indicates the exact number of atoms of each element contained in the smallest unit of a substance • Empirical Formula: indicates the lowest whole number ratio of atoms present in a substance • e.g.; hydrogen peroxide is a molecule that contains two atoms of oxygen and two atoms of hydrogen. • The molecular formula is H2O2 • The empirical formula is HO • For many molecules, the molecular and empirical formulas are identical (e.g.; H2O)

17. Chemical Formulas • For an ionic compound, there are no distinct molecules, so we use their empirical formula to identify them. • In the figure, each Na+is associated with six Cl-, each Cl- is associated with six Na+, and the whole solid is electrically neutral. • Other ionic compounds may have different structures and different charges (as long as the crystal remains electrically neutral) • ex.; Al3+ and O2- forms Al2O3

18. Nomenclature of Inorganic Compounds: Ionic Compounds • For ionic compounds, we name the cation followed by the anion • The anion or the cation is sometimes a polyatomic ion • ex.; NH4+ : ammonium CO32- : carbonate OH- : hydroxide PO43- : phosphate NO3- : nitrate SO42- : sulfate NO2- : nitrite SO32- : sulfite • ex.; KBr : potassium bromide ZnI2 : zinc iodide Al2O3 : aluminium oxide NaOH : sodium hydroxide KCN : potassium cyanide NH4Cl : ammonium chloride

19. Nomenclature of Inorganic Compounds: Covalent Compounds • Covalent compounds consist of distinct molecules • For a binary covalent compound, we name the first element in the formula first • ex.; HCl : hydrogen chloride NO : nitric oxide SiC : silicon carbide

20. Nomenclature of Inorganic Compounds: Covalent Compounds • When two elements form several compounds, Greek prefixes are used to specify the number of atoms of each element • e.g.; CO : carbon monoxide CO2 : carbon dioxide NO2 : nitrogen dioxide N2O4 : dinitrogen tetroxide • Many covalent compounds containing hydrogen are referred to by their common name • e.g.; B2H6 : diborane CH4 : methane NH3 : ammonia H2O : water SiH4 : silane PH3 : phosphine

21. Nomenclature of Inorganic Compounds: Acids and Bases • A definition of an acid is: a substance that releases an H+ when dissolved in water. • If the acid is not an oxyacid, i.e., does not contain oxygen, we use the prefix “hydro” and the suffix “ic” • ex.; HF : hydrofluoric acid HI : hydroiodic acid H2S : hydrosulfuric acid • N.B. HCl(g) is hydrogen chloride, but once it is dissolved in water, it become hydrochloric acid and releases H+

22. Nomenclature of Inorganic Compounds: Acids and Bases • For an oxyacid (general formula HmXOn) it often happens that there are multiple possible values of n for each element X, and as such, within this series of compounds, • There is always an acid in the series that ends with “ic” • Adding another oxygen to the “ic” acid produces the “per….ic” acid • Removing an oxygen atom to the “ic” acid produces the “ous” acid • Removing two oxygen atoms to the “ic” acid produces the “hypo…ous” acid

23. Nomenclature of Inorganic Compounds: Acids and Bases • ex.; H2CO3 : carbonic acid HNO3 : nitric acid HNO2 : nitrous acid H2SO4 : sulfuric acid H2SO3 : sulfurous acid HClO4 : perchloric acid HClO3 : chloric acid HClO2 : chlorous acid HOCl : hypochlorous acid

24. Nomenclature of Inorganic Compounds: Acids and Bases • When an oxyacid loses its H+ to produce an oxyanion • The “ic” acids produce the “ate” anions e.g.; ClO4- : perchlorate ClO3- : chlorate • The “ous” acids produce the “ite” anions e.g.; ClO2- : chlorite OCl- : hypochlorite • When not all of the possible H+ ions are lost, we have to specify the number of H’s present e.g.; H3PO4 : phosphoric acid H2PO4- : dihydrogenphosphate HPO42- : hydrogenphosphate PO43- : phosphate