Chapter 1 MATTER 1.1 Atoms and Molecules
At the end of this topic, students should be able to: • Identify and 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? • Matteris 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.
Protons and neutrons are found in the nucleus of the atom, while electrons surround the nucleus to form ‘electron cloud’.
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 Atom or ion Proton number
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.
Example: Number of protons = proton number = 13 Number of neutrons = nucleon number – proton number = 27 - 13 = 14 Number of electron = proton number – charge carried by species = 13 – ( +3) = 10
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 207. • Write out the atomic notation for lead. • 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 1214 66 1 1 2 1 3 1 H H H C C ;
Isotopes of an element have the same: • number of protons (proton number) • charge of nucleus of the atoms (ionization energy; electron affinity; size of the atom; electronegativity are the same) • number of electrons in a neutral atom • electronic configuration (the number of valence electrons) • chemical properties
Isotopes of an element have different: • number of neutrons (nucleon number) in the nucleus of the atoms • relative isotopic mass • 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: H2, Cl2, HCl, CO) • Polyatomic molecule: contains more than two atoms (example: H2O, NH3)
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: Mg2+, Al3+) • Anion: a negatively charged ion (number e > number p) (example: Cl, OH)
Monatomicion: ion contains only one nucleus (example: Fe3+, S2) • Polyatomic ion: ion contains more than one nucleus (example: H3O+, 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.
Ar = average mass of one atom of the element 1/12 x mass of one atom of C-12 • The mass of C-12 is defined exactly 12.00 a.m.u. • 1 atomic mass unit : a mass unit equal to 1/12 the mass of a C-12 atom
Example: Oxygen consists of three isotopes 16O ;17O and 18O in the ratio of 99.76 : 0.04 : 0.20. Calculate the relative atomic mass (to 4 decimal point) of oxygen. ANSWER : Relative Atomic Mass = (16 x 99.76) + (17 x 0.04) + (18 x 0.20) (99.76 + 0.04 + 0.20) = 16.0044
Exercise: Determine the Ar for Ne consists of three isotopes 20Ne ; 21Ne and 22Ne in the ratio of 90.92 ; 0.26 and 8.82. Answer: 20.18
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, CO2. Answer: Mr CO2 = Ar C + 2 Ar O = 12 + 2(16) = 44
Exercise: The relative molecular mass of ethanol, C2H5OH Answer: 46.00
1.1.7 Mass Spectrometer Mass spectrometer is used to determine: • Relative atomic mass of an element • Relative molecular mass of a compound • Types of isotopes that are found in the naturally occurring element including the abundance of the isotopes and its relative isotopic mass. • Recognize the structure of the compound in an unknown sample
A B 37Cl+ Vacuum 35Cl+ Vaporisation chamber Acceleration chamber (Electric field) Magnetic field Ionization chamber Ion detector There are five main stages:
(a) Vaporisation Chamber ~Sample of the element is vaporised into gaseous atom (b) Ionisation Chamber ~A gaseous sample (atom or molecule) is bombarded by a stream of high-energy electrons that are emitted from a hot filament. Collisions between the electrons and the gaseous atom (or molecule) produce positive ions by dislodging an electron from each atom or molecule. M sample + efilament M+ sample + e sample + efilament M M+
(c) Acceleration Chamber (Electric field) • The positive ions are accelerated by an electric field towards the two oppositely charged plates. The electric field is produced by a high voltage between the two plates. The emerging ions are of high and constant velocity.
Magnetic field • The positive ions are separated and deflected into a circular path by a magnet according to its mass/charge (m/e) ratio. • Positive ions with small m/eratio are deflected most and appear near A. Ions with large m/e ratio are deflected least and appear near B (Figure 1.1.2).
(e) Ion detector • The numbers of ions and types of isotopes are recorded as a mass spectrum. • In practice, the ion detector is kept in a fixed position. The magnetic field is varied so that the positive ions of different masses arrive at the detector at different times.
Mass spectrum • the horizontal axis – • the m/e ratio • nucleon number • isotopic mass • 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.
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 • Thus, the relative atomic mass of the element can be determined
Relative abundance 63 9.1 8.1 0 24 25 26 m/e The mass spectrum of magnesium shows that naturally occurring magnesium consists of three isotopes: 24Mg, 25Mg and 26Mg.
The height of each line is proportional to the abundance of each isotope. • In this example, 24Mg is the most abundance of the three isotopes. Ar of Mg: = 24.33
Average atomic mass = 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
Notes : • The height of each peak measures the relative abundance of the ion which gives rise to that peak. • 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
1.1.8 IUPAC Nomenclature for Cations, Anions and Salts • Cations • A positive ion is formed by the loss of a certain number of electrons from the given atom • For example: • Mg 2e- + Mg2+ (2e- were lost by Mg atom) • K 1e- + K+(1e- were lost by K atom)
Anions • A negative ion is formed by gain of one or more electrons by the given atom • For example: • S + 2e- S2-(2e- were gain by sulphur atom) • Cl + 1e- Cl-(1e- were gain by chlorine atom)
Salts - Combination of cation and anion to form ionic compound • For example: • Na+ + Cl- NaCl • Mg2+ + 2Br- MgBr2
a)Naming monoatomic ions • Monoatomic ions:ions consisting of single atomic nucleus • Cation: named by just adding ‘ion’ Ex: K+ : potassium ion Ca2+ : calcium ion • Anion : named by adding suffix ‘ide’ Ex: Cl- : chloride ion S2- : sulphide ion O2- : Oxide ion
b) Naming ionic compounds When naming an ionic compound made up of two element (binary): • the name of the cation (metal ion) appear first, followed by the name of the anion(non-metal ion)
System for naming cation of transition metals • Stock system/systematic system: - a Roman numeral that match the ionic charge is placed in bracket immediately after the element name of metal - for example: Fe 2+ iron(II) ion
Common nomenclature system - suffix-ous indicates the lower ionic charge and the suffix-ic indicates the higher ionic charges - for example: Fe2+ ferrous ion Fe3+ ferric ion
Rules for writing for an ionic compound • The positive ion is always written first. • The ratio of positive ions to negative ions must be such that the total number of +ve charge equal the total number of -ve charge; the formula unit must be electrically neutral • The smallest set of subscripts that give electrical neutrality is always chosen. We always write empirical formulas for ionic compounds
c)Naming polyatomic ions • Polyatomic ions are composed of 2 or more atoms bonded together • Most polyatomic ion consist to non-metal such as P, S, C, or N bonded to oxygen atoms • -ate: most common polyatomic ion • -ite : used for names of related ion that have one less oxygen atom
Chapter 1 MATTER 1.2 Mole Concept