Chapter 1

1.1 Atoms and Molecules At the end of this topic, students should be able to: Describe proton, electron and neutron. Define proton no.,Z, nucleon no., A and isotope. Write isotope notation. Define relative atomic mass and relative molecular mass based on the C-12 scale. Sketch and explain the function of the following main components of a simple mass spectrum: Analyze mass spectrum of an element. Name cations, anions and salt according to the IUPAC.

1.1 Atoms and Molecules What is matter? Matter is anything that has mass and occupies space. All matters consist of tiny particles called atoms. 3 states of matter are solid, liquid and gas. For examples are metal, plastics, gas, etc.

1.1.1 Atoms

Protons and neutrons are found in the nucleus of the atom, while electrons surround the nucleus to form ‘electron cloud’.

The properties of these particles are summarized in the table.

Nucleon the particles that are found in the nucleus. consist proton and neutron. By proton we can identified the element.

PROTON NUMBER AND NUCLEON NUMBER Proton number (atomic number), Z.  The proton number = the number of protons in the nucleus of an atom. Nucleon number (mass number), A.  The nucleon number = the total number of protons and neutrons in the nucleus of an atom.

1.1.2 Isotopic notation Nucleon number Proton number Atom or ion

For neutral atom, number of protons = number of electrons. For +ve ion, number of electron is less than number of protons. For –ve ion, number of electrons is more than number of protons.

Number of protons = proton number = 13 Number of neutrons = nucleon number – proton number = = 14 Number of electron= proton number – charge carried by species = 13 – ( +3) = 10 Example:

Exercise: Determine the number of proton, neutron and electron in the following species. a) b)

Exercise: The atomic number of lead (Pb) is 82 and the mass number is Write out the atomic notation for lead. 2.Give the number of protons, neutrons and electrons.

1.1.3 ISOTOPES Two or more atoms of the same element having same proton number but different nucleon number.

Eg : 1) hydrogen isotope 2) carbon isotope H 2121 H 3131 H C C ;

Isotopes of an element have the same: 1.number of protons (proton number) 2.charge of nucleus of the atoms (ionization energy; electron affinity; size of the atom; electronegativity are the same) 3.number of electrons in a neutral atom 4.electronic configuration (the number of valence electrons) 5.chemical properties

Isotopes of an element have different: 1.number of neutrons (nucleon number) in the nucleus of the atoms 2.relative isotopic mass 3.physical properties (e.g boiling point / melting point, density, effusion rate,…)

1.1.4 Molecule A molecule consists of a small number of atoms joined together by covalent bond. Diatomic molecule: contains two atoms (example: H 2, Cl 2, HCl, CO) Polyatomic molecule: contains more than two atoms (example: H 2 O, NH 3 )

1.1.5 Ion An ion is a charged species formed from a neutral atom or molecule when electrons are gained or lost as the result of a chemical reaction. Cation: a positively charged ion (number e < number p) (example: Mg 2+, Al 3+ ) Anion: a negatively charged ion (number e > number p) (example: Cl , OH  )

Monatomic ion: ion contains only one nucleus (example: Fe 3+, S 2  ) Polyatomic ion: ion contains more than one nucleus (example: H 3 O +, CN  )

1.1.6 RELATIVE MASS Isotopes carbon -12 as a reference or standard for comparing the masses of other atoms. RELATIVE ATOMIC MASS, Ar Average mass of one atom of the element relative to 1/12 times the mass of one atom of carbon-12.

The mass of C-12 is defined exactly a.m.u. 1 atomic mass unit : a mass unit equal to 1/12 the mass of a C-12 atom Ar = average mass of one atom of the element 1/12 x mass of one atom of C-12

Example: Oxygen consists of three isotopes 16 O ; 17 O and 18 O in the ratio of : 0.04 : Calculate the relative atomic mass (to 4 decimal point) of oxygen. ANSWER :

RELATIVE MOLECULAR MASS, Mr The mass of one molecule of the substance relative to 1/12 times the mass of one atom of carbon-12. = sum of the relative atomic masses of all the atoms shown in the molecular formula. Mr = average mass of one molecule of the substance 1/12 x mass of one atom of C-12

Example: The relative molecular mass of carbon dioxide, CO 2. Answer: Mr CO 2 = Ar C + 2 Ar O = (16) = 44

Mass spectrum the horizontal axis – i. the m/e ii. nucleon iii. isotopic iv. relative atomic mass of the ions entering the detector. The vertical axis - the abundance or detector current or relative abundance or ion intensity or percentage abundance of the ions.

1.Information from a mass spectrum of an element the isotopes which are present in the element the relative isotopic mass of each isotope the abundance of each isotope 2.Thus, the relative atomic mass of the element can be determined

Relative abundance m/e The mass spectrum of magnesium shows that naturally occurring magnesium consists of three isotopes: 24 Mg, 25 Mg and 26 Mg.

The height of each line is proportional to the abundance of each isotope. In this example, 24 Mg is the most abundance of the three isotopes. A r of Mg: = 24.33

where Q = the abundance of an isotope of the element = the percentage of the isotope found in the naturally occurring element m = the relative isotopic mass of the element Average atomic mass =

Notes : 1)The height of each peak measures the relative abundance of the ion which gives rise to that peak. 2)The total number of peaks in the mass spectrum of an element shows the types of naturally occurring isotopes. 3)The ratio of mass/charge for each species is found from the value of the accelerating voltage associated with a particular peak. Many ions have a charge of +1 elementary charge unit, and the ratio m/e is numerically equal to m, the mass of the ion. (1 elementary charge unit = 1.60x10  19 C) 4)The ion with the highest value of m/e is the molecular ion, and its mass gives the molecular mass of the compound