Understanding the Empirical Formula of Copper Oxide

Copper forms two oxides, Cu2O and CuO, each with its own empirical formula. Understanding these oxides is crucial for many applications from industrial uses to chemical analysis.

Types of Copper Oxides

Copper can form two distinct oxides depending on the valence state of copper:

Cuprous Oxide (Cu2O): A red-coloured solid used in certain antifouling paints. Cupric Oxide (CuO): A black solid and an ionic compound.

Empirical Formula Calculation

To determine the empirical formula of copper oxide, the following steps are typically followed:

Understand the atomic masses: ( Cu 63.5 , g/mol ) and ( O 16 , g/mol ). For the given masses, divide the mass of each element by its respective molar mass. Cu: ( 127 / 63.5 2 ) O: ( 32 / 16 2 ) Divide by the least common factor, which is 2 in this case: Cu: ( 2 / 2 1 ) O: ( 2 / 2 1 ) The empirical formula for copper oxide is therefore CuO.

Chemical Equation and Reaction

The empirical formula can also be derived using the chemical equation for the formation of copper oxide. For instance:

Reaction: ( 2Cu O_2 rightarrow 2CuO )

Consequently, the empirical formula of copper oxide is CuO.

Empirical Formula from Sample Analysis

Experimental data from a copper oxide sample containing 0.150 g Cu and 0.038 g O can be used to determine its empirical formula. Calculations reveal the following:

0.150 g Cu corresponds to 0.00236 mol Cu 0.038 g O corresponds to 0.00238 mol O The ratio of Cu to O is 1:1, indicating an empirical formula of CuO. This is a reasonable answer considering copper usually enters its compounds in the 2 oxidation state.

Crystal Structure of Copper Oxide

The compound is ionic, and its empirical formula (CuO) is also its molecular formula. Copper(II) oxide is composed of crystals containing equal numbers of Cu2 and O2- ions.

Additional Data

In another scenario, to determine the empirical formula of a sample, the following steps can be used:

Mass of copper oxide: 10 g Mass of copper: 8.88 g Mass of oxygen: 10 g - 8.88 g 1.12 g Moles of Cu: ( 8.88 / 63.5 0.136 , mol ) Moles of O: ( 1.12 / 16.0 0.070 , mol ) Divide by the smallest number of moles (0.070): Cu: ( 0.136 / 0.070 2 ) O: ( 0.070 / 0.070 1 ) This confirms the empirical formula Cu2O (copper(I) oxide), a red powder.

Understanding the empirical formula of copper oxides is not only important for chemical analysis but also for various industrial applications, including the production of certain resistant coatings and pigments.