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Preliminary Course Atomic Structure 1 + 2 Colm Healy Outline Part One • Matter • Atomic Structure • Atomic Model(s) Part Two • Periodic Table • Main Group Properties • Ions 2 Chemistry The Study of Stuff “Matter” – physical substances Properties depend on underlying with mass and volume molecular or atomic structure 3 The Basics All Matter is made up of Atoms Tiny (~1x10-10 m) particles Ultimately the source of all properties of matter 4 The Basics There are 118 different types of atoms – these are called the Elements. 5 The Basics Atoms bind together to form Molecules. Single atoms Complicated molecules Very Complicated molecules Molecules are discrete entities with discrete properties Atoms form molecules in a very predictable way, based on their elements 6 The Basics If a substance contains more than one element chemically bound together, this is called a Compound If a substance contains multiple elements or compounds that are not bound together, it is called a MIXTURE. 7 To Recap: Classifications of Matter 8 In more detail Atomic Theory How exactly does atomic structure dictate properties? ~500BC, Democritus hypothesises atoms are different shapes and sizes Atoms can recombine in different ways 9 In more detail Atomic Theory John Dalton (1803) updates the theory: • Matter is made of extremely small particles called atoms. • Atoms of different elements differ in size, mass, and other properties. • Atoms cannot be subdivided, created, or destroyed. • Atoms of different elements combine in simple whole-number ratios to form chemical compounds. • In chemical reactions, atoms are combined, separated, or rearranged. 10 Atomic Theory Experimental Evidence The Law of Multiple Proportions: Combinations of atoms always occur in whole number ratios: Examples: Water H2O Peroxide H2O2 Carbon dioxide CO2 Carbon monoxide CO Ammonia NH3 11 Atomic Theory Charged particles Dalton’s atomic theory didn’t explain WHY atoms were combining into compounds Around 1900 people discover electrically charged particles even smaller than atoms JJ Tomson’s experiments (1897) 12 Electromagnetic particles Thompson and Rutherford introduce ideas of the electron and the nucleus: Electrons are NEGATIVELY charged Nucleus is POSITIVELY charged Electrons orbit the nucleus – the atom is mostly empty space Rutherford’s experiments 1910 13 Some Basic Physics With (electrically) charged particles, OPPOSITES ATTRACT Coulomb’s Law Conversely, like charges repel Atomic Attraction 14 Atomic Theory The last piece of the puzzle Chadwick divides the nucleus into Protons and Neutrons The nucleus is very small : Atom approx. 10-10 m Nucleus approx. 10-14 m But most of the mass (>99%) is in the nucleus 15 Atomic Theory Particles Recap *1 amu = 1.66053904 × 10-24 grams Name(Symbol) Mass (amu*) Charge Proton (p+) 1.00727 1+ Neutron (n0) 1.00866 0 Electron (e-) 0.00054858 1- Nucleus Atoms are generally neutral ie. Number of protons = number of electrons There can be different numbers of neutrons – this gives rise to ISOTOPES 16 Atomic Structure There are more details (to do with the location of the electrons and the amount of energy they can have), but these will be dealt with in more detail later 17 The Periodic Table Lists all elements (atom types) Very useful tool to determine properties of elements 18 The Periodic Table Columns are called GROUPS. Elements in a group all have similar properties. Rows are called PERIODS. Properties change from period to period in a predictable way. 19 The Periodic Table How to read it: Atomic number (Z). Number of protons in the nucleus Always a whole number Number of protons = number of electrons Mass number (A) Mass/weight of an atom in amu Need this to calculate masses etc. 20 Isotopes Atomic number (Z) = #protons Mass number (A) ≈ #protons + #neutrons Therefore #neutrons = A - Z ISOTOPES have same Z but different A (ie. Different numbers of neutrons) A on periodic table is average of all naturally occurring isotopes, so not necessarily a whole number 21 The periodic table A tool for predicting properties The periodic table is arranged by increasing Z (atomic number) IE. The different elements are defined by the number of protons Adding protons also adds electrons Properties “repeat” because electrons shells get filled up 22 Group 1 – Alkali Metals • Soft, silvery metals • Low melting points • Very reactive with water or air • Produce H2 when reacted with water • Increasing reactivity down the group 23 Group 2 – Alkali Earth Metals • Harder, silvery metals • Produce H2 when reacted with water • Much less reactive • Increasing reactivity down the group 24 Group 3-11 – The transition metals • Various properties • Quite reactive (good catalysts) • Very important in chemistry and biology • Examples include iron, copper, gold, platinum, mercury…. 25 Group 17 - Halogens • Very reactive and toxic • Used as disinfectants • Form acids with hydrogen (HCl, HBr) • Decreasing reactivity down the group 26 Group 18 – The Noble Gases • Colourless, odourless gases • Almost completely inert • Almost never react • Electronically Stable 27 The Noble Gases A special case The noble gases are the most stable elements This is to do with their electron configuration (more details later) Because the noble gases appeared (roughly) every eight elements, this is sometimes called the OCTET RULE 28 The “octet rule” In other words, atoms want to obtain a “noble gas configuration” Atoms will try and find the shortest path to a noble gas configuration Want to LOSE ELECTRONS Want to GAIN ELECTRONS 29 Sodium Neon (noble gas) has 10 electrons (Z=10). It is unreactive. Sodium (group 1 metal) has 11 electrons (Z=11) Therefore sodium desperately wants to lose an electron to gain a noble gas configuration 30 Chlorine Similarly, chlorine (group 17 halogen) has 17 electrons (Z=17) The nearest noble gas is argon (Ar), which has 18 electrons (Z=18). Chlorine will therefore pull an electron from something to satisfy the octet rule 31 Forming Salt So if chlorine wants to gain an electron, and sodium wants to lose one, we can expect them to react together: Sodium Chlorine Sodium Chloride (table salt) Salts like this are made up of IONS (permanently charged atoms) The positive one (lost an electron) is called the CATION The negative one (gained an electron) is called the ANION Remember, opposites attract, so there is a force binding the ions together into a compound 32 Cations + Anions Metals tend to form CATIONS Want to LOSE ELECTRONS NON- Metals tend to form ANIONS Want to GAIN ELECTRONS 33 Multiple ions We can predict a lot about stable compounds from this For example, group 2 metals need to lose two electrons to obtain noble gas configuration Li and Na (Group 1) Li+ Na+ (Notice charge = group) Mg and Ca (Group 2) Mg2+ Ca2+ So Calcium would have to react with TWO clorines – Calcium forms CaCl2, not CaCl 34 Multiple Ions Similarly, with groups 15-17 (non-metals): Nitrogen (Group 15) N3- Oxygen (Group 16) O2- Chlorine (Group 17) Cl- (Notice charge = group - 18) Oxygen wants to gain two electrons – it will form Na2O or CaO 35 Working out compounds Add up positive charges (number of groups from the left) and subtract negative charges (number of groups from right) until things are neutral This method is not 100% foolproof, as different types of bonding come in to play (especially with carbon or transition metals), but this predicts a huge amount of compounds Examples: Water H2O Carbon dioxide CO2 Ammonia NH3 NaCl KBr HCl 36 Best of luck with first year!