# Significant Figures

Every measurement has some amount of uncertainity associated with it. A clear picture of this is given by precision (closeness of various measurements for the same quantity) and accuracy (agreement of a particular value to the true value. This problem of uncertainity was solved by considering the meaningful digits that are certain. They are called significant figures.

There are certain rules for determining the no of significant figures. They are listed below:
1. All non zero digits are significant. For example 0.359 has 3 significant figures.
2. Zeros preceding to first non zero digit are not significant (eg. 0.002355 has only 4 significant figures.
1. Q Zeros occurring between the non zero digits are significant (eg. 30.002 has five significant digits.
3. Zeros at the end or right of a no. are significant provided they are on the right side of a decimal point
(eg. 0.3000 has 4 significant figures where as 3000 has only 1 significant figure.
4. The above fact can be well understood by writing nos followed by zeros with scientific notation
(eg. 100 = 1 X 10 2 has 1 significant figure whereas 1.0 X 10 2 , 1.00 X 10 2 has 2 and 3 significant figures respectively.

# Reactions in solutions

Expressing strengths solutions:
1. Mass percent
Mass percent = Mass of solute/Mass of solution
2. Mole fraction
If A and B are the compositions of the two solutions respectively, then their molefractions are given by
XA = nA/nA + nB
XB = nB/nA+nB
3. Molarity(M)
Molarity = no of moles of the solute/ Volume of the solution in litres
4. Molality (m)
Molality = no of moles of the solute / Mass of the solvent in Kg

Limiting Reagent: The reactant required for a chemical reaction not present in sufficient quantity limits the the formation of the product. Such reactants which are consumed faster and not available for the completion of the reaction is known limiting reagent.