Naming compounds writing formulas and equations
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Naming Compounds Writing Formulas and Equations. Larry Scheffler Lincoln High School. Naming Compounds. The chemical formula represents the composition of each molecule.

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Naming Compounds Writing Formulas and Equations

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Naming compounds writing formulas and equations

Naming CompoundsWriting Formulasand Equations

Larry Scheffler

Lincoln High School


Naming compounds writing formulas and equations

Naming Compounds

The chemical formula represents the composition of each molecule.

In writing the chemical formula, in almost all cases the element farthest to the left of the periodic table is written first.

So for example the chemical formula of a compound that contains one sulfur atom and six fluorine atoms is SF6.

If the two elements are in the same period, the symbol of the element of that is lower in the group (i.e. heavier) is written first e.g. IF3.


Naming compounds writing formulas and equations

Naming Ionic Compounds

Ionic compounds are combinations of positive

and negative ions.

In writing the chemical formula the positive ion is

written first, It is then followed by the name of the

negative ion.

Monatomic anions end in ide. Special endings

apply for polyatomic ions

Examples

NaCl Sodium chloride

BaF2 Barium Fluoride

ZnO Zinc Oxide


Names of polyatomic ions with oxygen

Names of Polyatomic Ions with Oxygen

  • Polyatomic ions usually contain oxygen in addition to another element. 

  • Normally they have a negative charge. 

  • They end in either "ate" or "ite" depending on the number of oxygen atoms present.


Polyatomic ion exceptions

Polyatomic Ion -- Exceptions

  • Most polyatomic ions contain oxygen

  • Their names end in “ite” or “ate”.

  • There are several exceptions

    OH- hydroxide

    CN- cyanide

    SCN- thiocyanate


Elements with multiple cations

Elements with Multiple Cations

  • When an element can form more than one cation a Roman numeral is used to distinguish the oxidation state of the compound.

  • Iron, Tin, Lead, Copper, and are common elements with more than one cation.

  • Examples

  • PbSO4  =  lead (II) sulfate  This compound is formed from Pb2+ and  SO42-

  • Pb(SO4)2 =  lead (IV) sulfate  This compound is formed from Pb4+ and  SO42-

  • Fe(OH)2  =  iron (II) hydroxide  This compound is formed from Fe2+ and  OH-

  • Fe(OH)3  =  iron (III)  hydroxide  This compound is formed from Fe3+ and  OH-


Examples of ionic compounds

Examples of Ionic Compounds

  • NaCl=Sodium chloride

  • ZnF2=Zinc fluoride

  • KOH=Potassium hydroxide

  • Ca(NO3)2=Calcium nitrate

  • BaSO3=Barium Sulfite

  • Al2(SO4) 3 =Aluminum sulfate

  • Ca3(PO3)2=Calcium phosphite

  • NH4Cl=Ammonium chloride

  • (NH4)2CO3=Ammonium carbonate


Naming compounds writing formulas and equations

Naming Covalent Compounds

When naming covalent compounds, the name of the first element in the formula is unchanged.

The suffix “-ide” is added to the second element.

Often a prefix to the name of the second element indicates the number of the element in the compound

Examples:

SF6 – sulfur hexafluoride

P4O10 – tetraphosphorous decoxide

CO – carbon monoxide

CO2 – carbon dioxide


Covalent molecules with multiple possibilities

Covalent molecules with multiple possibilities

  • A Roman Numeral is used to indicate the state of the more positive element

  • Examples

  • N2O   =  Nitrogen (I) oxide  Since oxygen has a 2- charge, the nitrogen must be 1+ to  balance the charges.    Also known asdinitrogen monoxide

  • N2O3 =  Nitrogen (III) oxide   Since oxygen has a 2- charge, the nitrogen must be 3+ to balance the charges  Also  known as dinitrogen trioxide


Naming compounds writing formulas and equations

Binary compounds of Hydrogen

Water H2O is not called dihydrogen monoxide

The binary compounds of hydrogen are special cases. They were discovered before a convention was adopted and hence their original names have stayed.

  • Hydrogen forms binary compounds with almost all non-metals except the noble gases.

  • Examples

    • HF - hydrogen fluoride

    • HCl - hydrogen chloride

    • H2S - hydrogen sulfide


Naming compounds writing formulas and equations

Acids

When many hydrogen compounds are dissolve in water they take on the form of an acid. Special rules apply to acids. The “ite” suffix becomes “ous” and the “ate” suffix becomes “ic”


Writing formulas for ionic compounds

Writing Formulas for Ionic Compounds

  • Write the positive ion (cation) first, then the negative ion.

  • The positive charges must balance the negative charges.

  • Use subscripts to show how many times each ion must appear in order for the charges to balance. A subscript is not used if the ion appears only once

  • Use parenthesis around polyatomic ions that appear more than once in the formula


Examples

Examples

  • Na+ and Cl- = NaCl

  • Zn2+ and Br- = ZnBr2

  • K+ and OH- = KOH

  • Ca2+ and OH- = Ca(OH)2

  • Fe2+ and SO42- = FeSO4

  • Fe3+ and SO42- = Fe2(SO4) 3

  • Ca2 + and PO43- = Ca3(PO4)2

  • NH4+ and Cl- = NH4Cl

  • NH4+ and CO32- = (NH4)2CO3


Chemical reactions

Chemical Reactions

  • Elements and compounds frequently undergo chemical reactions to form new substances

  • In a chemical reaction, chemical bonds are frequently broken and new chemical bonds are formed

  • Atoms are neither created nor destroyed in an ordinary chemical change


Chemical reactions1

Chemical Reactions

  • A balanced chemical reaction is used to describe the process that occurs in a chemical change.

  • For example: Zinc reacts with hydrochloric acid to produce zinc chloride and hydrogen gas.

  • This chemical reaction could be written as

    Zn + 2 HCl  ZnCl2 + H2


Reactants and products

Reactants and Products

  • In the chemical reaction

    Zn + 2 HCl  ZnCl2 + H2

    Reactants Products

  • This shorthand way of describing a chemical reaction is known as a chemical equation

  • The starting materials are shown on the left and are known as reactants

  • The substances formed are shown on the right and are known as the products


Balancing a chemical reaction

Balancing a Chemical Reaction

  • A proper chemical reaction must be balanced

    Zn + 2 HCl  ZnCl2 + H2

    Reactants Products

  • Each element must appear on both sides of the arrow and equal number of times

  • Chemical reactions can be balanced by inserting numbers in front of formulas.

  • These numbers are called coefficients


Balancing chemical reactions

Balancing Chemical Reactions

  • Most simple equations can be balanced by inspection

  • Example: Balance the following equation

    BaCl2 + K3PO4 Ba3 (PO4)2 + KCl

  • There are 3 Ba on the right so we need coefficient of 3 in front of BaCl2

  • There are 2 PO4 on the right so we need a coefficient of 2 in front of K3PO4.

  • This leaves 6 K on the left so we need a coefficient of6 in front of the KCl on the right

    The balanced equation is

    3 BaCl2 + 2 K3PO4 Ba3 (PO4)2 + 6 KCl


Balancing chemical reactions1

Balancing Chemical Reactions

  • An equation is balanced when there are the same number and kind of atoms on both sides of the arrow

    3 BaCl2 + 2 K3PO4 Ba3(PO4)2 + 6 KCl


State symbols

State Symbols

  • State symbols are often added to chemical equations.

    CaCO3(s) + 2 HCl (aq) CaCl2(aq) + CO2(g) + H2O (l)


Types of reactions

Types of Reactions

  • There are many kinds of chemical reactions that occur. Some are very simple while others are very complex and may occur in multiple steps.

  • A number of reactions conform to some relatively simple patterns

  • Understanding and identifying these patterns can be helpful in predicting the products of similar reactions


Direct combination

Direct Combination

  • In a direct combination, two elements or compounds combine to form a more complicated product

  • Examples

    CaO + CO2 CaCO3

    2 H2 + O2  2 H2O

    FeCl2 + Cl2  FeCl3

    N2 + O2 2 NO


Decomposition

Decomposition

  • In a dcecomposition, a single compound is broken down into two or more simplier substances

  • Examples

    2 KClO3 2 KCl + 3 O2

    ZnCO3  ZnO + CO2

    Cu(OH)2 CuO + H2O


Single replacement

Single Replacement

  • In a single replacement, one substance (usually an element) takes the place of another in a compound

  • Examples

    Zn + H2SO4ZnSO4 + H2

    Cl2+ 2 KBr 2 KCl+ Br2

    Mg + CuCl2  MgCl2 + Cu


Double replacement

Double Replacement

  • In a double replacement, two substances exchange places in their respective compounds

  • Examples

    AgNO3 + NaCl AgCl + NaNO3

    3 CaCl2 + 2 K3PO4 Ca3(PO4)2 + 6KClBaCl2 + Na2SO4 BaSO4 + 2NaCl


Diatomic molecules

Diatomic Molecules

  • Certain elements exist as diatomic molecules in nature


Diatomic molecules1

Diatomic Molecules

  • Certain elements exist as diatomic molecules in nature


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