Chapter 1 Notes Atoms and Bonding

Section 1: Elements and Atoms A. Matter is defined as anything that has mass and takes up space. - Ancient Greeks believed that all matter was made up of four elements; earth, air, fire and water.

B .Elements are the simplest pure substances, and they cannot be broken down into any other substance. (the periodic table of elements) - Often called the building blocks of matter because all matter is composed of one element or a combination of two or more elements. C. Compound – a pure substance made of two or more elements that are chemically combined. Ex. Sodium Chloride (table salt) = NaCl

D. Mixture – two or more substances that are in the same place but are not chemically combined. Ex. Air, soil, wood, orange juice

E. Atoms 1. Democritus (430 B.C.) Proposed the idea of an atom. 2. Matter is formed of pieces that are “uncuttable” 3. Modern definition: a. An atom is the smallest particle of an element.

2. Atomic Theory and Models A scientific theory- a well-tested idea that explains and connects a wide range of observations Daltons atomic theory- All elements are composed of atoms. All atoms of the same element are exactly alike and have the same mass. Atoms of different elements are different and have different masses. An atom of one element cannot be changed into an atom of a different element. Atoms cannot be created or destroyed in any chemical change, only rearranged. Every compound is composed of atoms of different elements, combined in a specific ratio.

Found atoms contained negatively charged particles called electrons.   Thomson- Found atoms contained negatively charged particles called electrons. Aka, “muffin model”

Rutherford- gold foil experiment Inferred that an atoms positive charge must be clustered in a tiny region in its center called the nucleus The nucleus is made of one or more positively charged particles called protons

Bohr’s Model-electrons orbit the nucleus in certain orbits similar to the way planets orbit the sun.

Electron Cloud Model-electrons can be anywhere in a cloudlike region around the nucleus. The “cloud” symbolizes where an electron is likely to be found. An electron’s movement is related to its energy level- specific amount of energy it has. Electrons with more energy are farther from the nucleus. Modern Atomic Model-James Chadwick discovered the neutron, a neutral particle in the nucleus with nearly the same mass as a proton.

Section 2: Atoms, Bonding, and the Periodic Table

1. Valence electrons and Bonding a. valence electrons- electrons that are in the highest energy level of an atom and are held loosely by the nucleus -valence electrons determine the properties of the element and how the atom can bond with other elements

b. electron dot diagrams-symbol of an element surrounded by the number of dots that equals the number of valence electrons the element has.

c. Chemical bond- the force of attraction that holds two atoms together as a result of the rearrangement of electrons between them -atoms are less likely to react with they have 8 valence electrons. -atoms will gain or lose electrons to become more stable. -when atoms bond, electrons can be transferred from one atom to another or are shared. -the bonding of atoms causes a chemical reaction and makes a new substance.

2. The Periodic Table A. First put together in the 1800’s by a Russian Chemist –Dimitri Mendeleev B. Modern periodic table is used worldwide 1. Elements are organized into groups with the same number of valence electrons in their atoms. 2. Important information contained within each block

a. atomic number = number of protons   1) protons are positively charged subatomic particles 2) nucleus is found in the center of the atom. It is made up of protons and neutrons. 3) neutron – is a neutral particle – nearly the same mass as a proton. 4) electron – typically equals the number of protons in the Atom, except if the atom is an ion, it is a negative particle that orbits around the nucleus in different energy levels. 5). atomic mass = the number of protons and neutrons in the nucleus of an atom.

Calculations involving Subatomic Particles: atomic number = # of protons mass number = # of protons + # of neutrons neutral atom: # of protons = # of electrons charged ion: # of electrons = # of protons - +charge # of electrons = # of protons + - charge

3. Organization of the periodic table (see periodic table handout)   Groups (family)– a column (vertical) of elements that have the same number of valence electrons. Period – elements in the same horizontal row of the periodic table. Elements are classified into three major categories that include metals, nonmetals, and metalloids.

4. Families of the periodic table a. alkali metals-very reactive; Li, Na, K, Rb, Cs, Fr b. alkaline earth- reactive; Be, Mg, Ca, Sr, Ba, Ra c. halogens-“salt formers”; F, Cl, Br, I d. noble gases- all stable, few react; He, NE, Ar, Kr, Xe, Rn

Section 3 Ionic Bonds

1. Ions and ionic bonds a. ion- an atom or group of atoms with a charge. -cation- an atom that loses an electrons and has a positive charge. -anion- an atom that gains an electrons and has a negative charge. b. polyatomic ion- ions that are made of more than one atom and can be positive or negative. c. ionic bonds-attraction between oppositely charged ions. d. an ionic compound is one that consists of positive and negative ions. e. cations are usually metals and anions are usually nonmetals

2. Chemical formulas and names Formula writing Ex: sodium chloride NaCl Magnesium chloride  MgCl2  

Formula units have a net zero charge Determining Charge Formula Unit-Simplest ratio of the ions represented in an ionic compound. Formula units have a net zero charge Determining Charge monoatomic ions –a one atom ion Charge based on the location of the element on the periodic table 3. oxidation number –charge of the monatomic ion 1A 1+ 6A 2- 2A 2+ 7A 1- 5A 3- 8A

Oxidation numbers

Determining the Formula : Criss-cross method Cations listed first and anions listed second Find the symbols and charges in the upper right corner of the symbol 3. Cross the charges, drop the (+) or (-) sign 4. Reduce if necessary and never write the number 1

Examples: potassium oxide ______________________ aluminum sulfide ______________________ calcium fluoride _______________________ zinc fluoride ________________________ aluminum iodide _______________________

Polyatomic ions (see handout for list) 1. Write the symbols and charges 2. Cross the charges and drop the sign 3. If the oxidation number going to the polyatomic is greater than 1, add parenthesis. 4. Reduce the charges if possible, but do not change the polyatomic

Naming Ionic Compounds (continued) II. Polyatomic ions each have specific names which must be memorized so they can be recognized on sight. (At this point, if you are asked to name any compound that contains more than two elements, it will contain at least one polyatomic ion.) A few of the more common polyatomic ions Formula Name C2H3O21- acetate CO32- carbonate HCO31- bicarbonate NH41+ ammonium Formula Name NO31- nitrate OH1- hydroxide PO43- phosphate SO42- sulfate

Examples: sodium nitrate ______________________ potassium acetate ___________________ aluminum sulfate_____________________ lithium chlorate ______________________

Naming ions and ionic compounds 1. oxyanions –polyatomic ion composed of an element, usually a metal and oxygen The ion with more oxygen atoms is named using the root of the nonmetal plus the suffix -ate. The ion with fewer oxygen atoms is named using the root of the nonmetal plus the suffix -ite

Rules for naming. 1)Name the cation first. 2)Monoatomic cations use the element name 3)Monoatomic anions take their name from the root of the element name and add –ide 4)Some metals can have multiple oxidation numbers – distinguish them by using Roman numerals Ex. Fe+2 and Fe+3 iron (II) iron (III) 5)Group I metals always +1 and Group II metals always +2 6)If the compound contains a polyatomic ion, simply name the ion.

Naming Ionic Compounds (continued) If a Roman numeral is required, the charge on the metal ion must be determined from the charge on the negative ion. Helpful Rules to Remember A metal ion is always positive. The Roman numeral indicates the charge, not the subscript. The positive and negative charges must cancel (total charge must = 0). Nonmetals are always negative & can never form more than one monatomic ion. Examples Formula Reasoning Name FeCl2 Cl has a 1- charge, and there are 2 of them for a total of 2-, so the Fe must be 2+ iron (II) chloride Fe2O3 O has a 2- charge, and there are 3 of them for a total of 6-, so the Fe must have a total charge of 6+ split equally between the two iron atoms, so each must have a 3+ charge iron (III) oxide PbS2 S has a 2- charge, and there are 2 of them for a total of 4-, so the Pb must be 4+ lead (IV) sulfide Cu3N N has a 3- charge, so the Cu must have a total charge of 3+ split equally between the 3 copper atoms, so each must have a 1+ charge copper (I) nitride

Additional Examples: C2H3O21- CO32- HCO31- NH41+ Na2SO4 Fe(NO3)2 AlCl3 PbI4 (NH4)3PO4 Mg3N2 C2H3O21- acetate CO32- carbonate HCO31- bicarbonate NH41+ ammonium NO31- nitrate OH1- hydroxide PO43- phosphate SO42- sulfate * Groups I & II, Al, Zn, Cd, and Ag need no Roman numeral.

AgC2H3O2 ________________ NaOH ___________________ NaClO3 ___________________ NaC2H3O2 _________________ BaSO4 ___________________ ZnCO3 ___________________ C2H3O21- acetate CO32- carbonate HCO31- bicarbonate NH41+ ammonium NO31- nitrate OH1- hydroxide PO43- phosphate SO42- sulfate * Groups I & II, Al, Zn, Cd, and Ag need no Roman numeral.

Naming Ionic Compounds: Examples C2H3O21- acetate CO32- carbonate HCO31- bicarbonate NH41+ ammonium NO31- nitrate OH1- hydroxide PO43- phosphate SO42- sulfate * Groups I & II, Al, Zn, Cd, and Ag need no Roman numeral. Na2SO4 sodium sulfate Fe(NO3)2 iron (II) nitrate AlCl3 aluminum chloride PbI4 lead (IV) iodide (NH4)3PO4 ammonium phosphate Mg3N2 magnesium nitride AgC2H3O2 silver acetate

Properties of Ionic Compounds a. hard, brittle crystals b. have high melting points c. solids at room temperature d. conduct electricity when dissolved in water

A covalent bond is one that results from the sharing of valence electrons A molecule is formed when two or more atoms bond covalently The difference between an ionic compound and a molecule is that an ionic compounds donate electrons and molecules share valence electrons 9.1 Covalent Bonds

Section 4 Covalent BOnds

How Covalent bonds form   How Covalent bonds form When two atoms share electrons it is called a covalent bond Covalent bonds form between nonmetals A molecule is a neutral group of atoms formed by covalent bonds Molecules can have single, double or triple bonds

2. Naming Molecular Compounds 1. The first element in the formula is always named first, using the entire element name. 2. The second element in the formula is named using the root of the element and adding the – ide 3. Prefixes are used to indicate the number of atoms of each type that are present (Exception: never use a prefix if there is only one of the first element)

Naming Covalent Compounds Prefixes Subscript Prefix 1 mono- 2 di- 3 tri- 4 tetra- 5 penta- Subscript Prefix 6 hexa- 7 hepta- 8 octa- 9 nona- 10 deca- Note: When a prefix ending in ‘o’ or ‘a’ is added to ‘oxide’, the final vowel in the prefix is dropped.

3. Writing Molecular Formulas A) Write the symbols for the elements B) Place the number of the prefix after the element symbol C) Do not cross charges or reduce

Examples: N2S4 NI3 XeF6 CCl4 P2O5 SO3 1 mono 2 di 3 tri 4 tetra 5 penta 6 hexa 7 hepta 8 octa 9 nona 10 deca * Second element in ‘ide’ from * Drop –a & -o before ‘oxide’ N2S4 NI3 XeF6 CCl4 P2O5 SO3

Naming Binary Covalent Compounds: Examples 1 mono 2 di 3 tri 4 tetra 5 penta 6 hexa 7 hepta 8 octa 9 nona 10 deca * Second element in ‘ide’ from * Drop –a & -o before ‘oxide’ N2S4 dinitrogen tetrasulfide NI3 nitrogen triiodide XeF6 xenon hexafluoride CCl4 carbon tetrachloride P2O5 diphosphorus pentoxide SO3 sulfur trioxide

Examples: 1 mono 2 di 3 tri 4 tetra 5 penta 6 hexa 7 heptaa 8 octa 9 nona 10 deca * Second element in ‘ide’ from * Drop –a & -o before ‘oxide’ xenon trioxide _____________________________ bromine trichloride _____________________________ nitrogen trifluoride _____________________________ tetrasulfur tetranitride _____________________________ ditellurium dichloride _____________________________ dinitrogen tetraoxide _____________________________ phosphorus pentachloride____________________________

xenon trioxide XeO3 bromine trichloride BrCl3 nitrogen trifluoride NF3 tetrasulfur tetranitride S4N4 ditellurium dichloride Te2Cl2 dinitrogen tetraoxide N2O4 phosphorus pentachloride PCl5

b. have low melting and boiling points c. most are gases and liquids 3. Properties of Molecular compounds a. molecular compound- compound of molecules that are covalently bonded b. have low melting and boiling points c. most are gases and liquids d. do not conduct electricity